• Shuffle
    Toggle On
    Toggle Off
  • Alphabetize
    Toggle On
    Toggle Off
  • Front First
    Toggle On
    Toggle Off
  • Both Sides
    Toggle On
    Toggle Off
  • Read
    Toggle On
    Toggle Off
Reading...
Front

Card Range To Study

through

image

Play button

image

Play button

image

Progress

1/304

Click to flip

Use LEFT and RIGHT arrow keys to navigate between flashcards;

Use UP and DOWN arrow keys to flip the card;

H to show hint;

A reads text to speech;

304 Cards in this Set

  • Front
  • Back
Abbreviations for the chemical elements.
1. Element Symbols
A given compound always contains elements in exactly the same proportion by mass.
2. Law of Constant Composition
A theory established by John Dalton in the early 1800’s, used to explain the nature of materials.
3. Dalton’s Atomic Theory
The fundamental unit of which elements are composed.
4. Atom
A substance with constant composition that can be broken down into elements by chemical processes.
5. Compound
A representation of a molecule in which the symbols for the elements are used to indicate the types of atoms present and subscripts are used to show the relative number of atoms.
6. Chemical Formula
A negatively charged particle that occupies the space around the nucleus of an atom.
7. Electron
The modern concept of the atom as having a dense center of positive charge (the nucleus) and electrons moving around the outside.
8. Nuclear Atom
The small dense center of positive charge in an ion.
9. Nucleus
A positively charged particle in an atomic nucleus.
10. Proton
The small dense center of positive charge in an ion.
11. Neutron
Atoms of the same element (the same number of protons) that have different numbers of neutrons; they have identical atomic numbers but different mass numbers.
12. Isotopes
The number of protons in the nucleus of an atom, each element has a unique ______ ______.
13. Atomic Number
The total number of protons and neutrons in the atomic nucleus of an atom.
14. Mass Number
A chart showing all the elements arranged in columns in such a way that all the elements in a given column exhibit similar chemical properties.
15. Periodic Table
A vertical column in the periodic table.
16. Groups
A group 1 metal.
17. Alkali Metals
A group 2 metal.
18. Alkaline Earth Metals
A group 7 element.
19. Halogens
A group 8 element.
20. Noble Gases
Several series of elements in which inner orbitals (d and f orbitals) are being filled.
21. Transition Metals
An element that gives up electrons relatively easily and is typically lustrous, malleable, and a good conductor of heat and electricity.
22. Metals
An element that does not exhibit metallic characteristics; chemically, a typical ________ accepts electrons from a metal.
23. Nonmetals
An element that has both metallic and nonmetallic properties.
24. Metalloids (Semimetals)
A molecule composed of two atoms.
25. Diatomic Molecule
An atom or a group of atoms that has a net positive or negative charge.
26. Ion
A positive ion.
27. Cation
A negative ion.
28. Anion
A compound that results when a metal reacts with a nonmetal to form cations and anions.
29. Ionic Compound
1. The Elements

• ___ known: 88 found in nature, others are man-made.
• Just as you had to learn the 26 letters of the alphabet before you learned to read and write, you need to learn the _____ and _______ of the chemical elements before you can read and write chemistry.
The Elements

• (115) known: 88 found in nature, others are man-made.
• Just as you had to learn the 26 letters of the alphabet before you learned to read and write, you need to learn the (NAMES) and (SYMBOLS) of the chemical elements before you can read and write chemistry.
2.	The Elements
2. The Elements
The Elements
The Elements
3. The Elements

How the Term Element is Used

• Could mean a ______ atom of that element (Ar or H).
• Could mean _________ of an element (H₂), which is hydrogen found in its natural state.
• Could mean _____ of elements are present in some form (sodium found in the human body).
• Look at each particular case to determine its proper use.
The Elements

How the Term Element is Used

• Could mean a (SINGLE) atom of that element (Ar or H).
• Could mean (MOLECULES) of an element (H₂), which is hydrogen found in its natural state.
• Could mean (ATOMS) of elements are present in some form (sodium found in the human body).
• Look at each particular case to determine its proper use.
4. Symbols for the Elements

• Each element has a ______ one- or two-letter symbol.
• First letter is always ___________ and the second is not.
• The symbol usually consists of the first one or two letters of the element’s name.
 Examples:
Oxygen = O
Krypton = Kr
• Sometimes the symbol is taken from the element’s original _____ or _____ name.
 Examples:
Gold = Au – aurum
Lead = Pb - plumbum
Symbols for the Elements

• Each element has a (UNIQUE) one- or two-letter symbol.
• First letter is always (CAPITALIZED) and the second is not.
• The symbol usually consists of the first one or two letters of the element’s name.
 Examples:
Oxygen = O
Krypton = Kr
• Sometimes the symbol is taken from the element’s original (LATIN) or (GREEK) name.
 Examples:
Gold = Au – aurum
Lead = Pb - plumbum
5.	Symbols for the Elements

Names and Symbols of the Most ______ Elements
5. Symbols for the Elements

Names and Symbols of the Most ______ Elements
Symbols for the Elements

Names and Symbols of the Most (COMMON) Elements
Symbols for the Elements

Names and Symbols of the Most (COMMON) Elements
6. Dalton’s Atomic Theory

1. Most natural materials are ________ of pure substances.
2. Pure substances are either elements or combinations of elements called _________.
3. A given compound always contains the same proportions (by mass) of the elements.
Dalton’s Atomic Theory

1. Most natural materials are (MIXTURES) of pure substances.
2. Pure substances are either elements or combinations of elements called (COMPOUNDS).
3. A given compound always contains the same proportions (by mass) of the elements.
7. Law of Constant Composition

• A given compound always has the same ___________, regardless of where it comes from.
 Water always contains 8 g of oxygen for every 1 g of hydrogen.
 Carbon dioxide always contains 2.7 g of oxygen for every 1 g of carbon.
Law of Constant Composition

• A given compound always has the same (COMPOSITION), regardless of where it comes from.
 Water always contains 8 g of oxygen for every 1 g of hydrogen.
 Carbon dioxide always contains 2.7 g of oxygen for every 1 g of carbon.
8. Dalton’s Atomic Theory (1808)

1. ________ are made of tiny particles called atoms.
2. All atoms of a given element are _________.
3. The atoms of a given element are _________ from those of any other element.
Dalton’s Atomic Theory (1808)

1. (ELEMENTS) are made of tiny particles called atoms.
2. All atoms of a given element are (IDENTICAL).
3. The atoms of a given element are (DIFFERENT) from those of any other element.
9.	Dalton’s Atomic Theory (Continued)

4.	Atoms of one element can combine with atoms of other elements to form _________. A given ________ always has the same ________ _______ and types of atoms.
9. Dalton’s Atomic Theory (Continued)

4. Atoms of one element can combine with atoms of other elements to form _________. A given ________ always has the same ________ _______ and types of atoms.
Dalton’s Atomic Theory (Continued)

4.	Atoms of one element can combine with atoms of other elements to form (COMPOUNDS). A given (COMPOUND) always has the same (RELATIVE NUMBERS) and types of atoms.
Dalton’s Atomic Theory (Continued)

4. Atoms of one element can combine with atoms of other elements to form (COMPOUNDS). A given (COMPOUND) always has the same (RELATIVE NUMBERS) and types of atoms.
10. Dalton’s Atomic Theory (Continued)

5. Atoms are indivisible in ________ _________. Atoms are not created or destroyed in chemical reactions. A chemical reaction simply changes the way the atoms are _______ together.
Dalton’s Atomic Theory (Continued)

5. Atoms are indivisible in (CHEMICAL PROCESSES). Atoms are not created or destroyed in chemical reactions. A chemical reaction simply changes the way the atoms are (GROUPED) together.
11. Dalton’s Atomic Theory

Concept Check

Which of the following statements regarding Dalton’s atomic theory are still believed to be true?

I. Elements are made of tiny particles called atoms.
II. All atoms of a given element are identical.
III. A given compound always has the same relative numbers and types of atoms.
IV. Atoms are indestructible.
Dalton’s Atomic Theory

Concept Check

Which of the following statements regarding Dalton’s atomic theory are still believed to be true?

I. Elements are made of tiny particles called atoms.
III. A given compound always has the same relative numbers and types of atoms.
12. Formulas of Compounds

Chemical Formulas Describe Compounds

• Compound – distinct substance that is composed of the atoms of two or more elements and always contains exactly the same ________ ______ of those elements.
• Chemical Formulas – expresses the types of atoms and the number of each type in each unit ([________)] of a given compound.
Formulas of Compounds

Chemical Formulas Describe Compounds

• Compound – distinct substance that is composed of the atoms of two or more elements and always contains exactly the same (RELATIVE MASSES) of those elements.
• Chemical Formulas – expresses the types of atoms and the number of each type in each unit ([MOLECULE]) of a given compound.
13.	Formulas of Compounds

Rules for Writing Formulas

1.	Each ____ present is represented by its element symbol.
2.	The number of each type of atom is indicated by a subscript written to the _____ of the element symbol.
3.	When only one atom of a g
13. Formulas of Compounds

Rules for Writing Formulas

1. Each ____ present is represented by its element symbol.
2. The number of each type of atom is indicated by a subscript written to the _____ of the element symbol.
3. When only one atom of a given type is present, the subscript 1 is not written.
Formulas of Compounds

Rules for Writing Formulas

1.	Each (ATOM) present is represented by its element symbol.
2.	The number of each type of atom is indicated by a subscript written to the (RIGHT) of the element symbol.
3.	When only one atom of a g
Formulas of Compounds

Rules for Writing Formulas

1. Each (ATOM) present is represented by its element symbol.
2. The number of each type of atom is indicated by a subscript written to the (RIGHT) of the element symbol.
3. When only one atom of a given type is present, the subscript 1 is not written.
14. Formulas of Compounds

Exercise

The pesticide known as DDT paralyzes insects by binding to their nerve cells, leading to uncontrolled firing of the nerves. Before most uses of DDT were banned in the U.S., many insects had developed a resistance to it. Write out the formula for DDT. It contains 14 carbon atoms, 9 hydrogen atoms, and 5 atoms of chlorine.
Formulas of Compounds

Exercise

The pesticide known as DDT paralyzes insects by binding to their nerve cells, leading to uncontrolled firing of the nerves. Before most uses of DDT were banned in the U.S., many insects had developed a resistance to it. Write out the formula for DDT. It contains 14 carbon atoms, 9 hydrogen atoms, and 5 atoms of chlorine.

C₁₄H₉Cl₅
15. The Structure of the Atom

J. J. Thomson (1898-1903)

• Postulated the existence of _________ using cathode-ray tubes.
• The atom must also contain ________ particles that balance exactly the ________ charge carried by particles that we now call electrons.
The Structure of the Atom

J. J. Thomson (1898-1903)

• Postulated the existence of (ELECTRONS) using cathode-ray tubes.
• The atom must also contain (POSITIVE) particles that balance exactly the (NEGATIVE) charge carried by particles that we now call electrons.
16.	The Structure of the Atom

_______-___ ____
16. The Structure of the Atom

_______-___ ____
The Structure of the Atom

(CATHODE-RAY TUBE)
The Structure of the Atom

(CATHODE-RAY TUBE)
17.	The Structure of the Atom

William Thomson (Plum Pudding Model)

•	Reasoned that the ____ might be thought of as a uniform “pudding” of positive charge with enough negative electrons scattered within to ______________ that positive charge.
17. The Structure of the Atom

William Thomson (Plum Pudding Model)

• Reasoned that the ____ might be thought of as a uniform “pudding” of positive charge with enough negative electrons scattered within to ______________ that positive charge.
The Structure of the Atom

William Thomson (Plum Pudding Model)

•	Reasoned that the (ATOM) might be thought of as a uniform “pudding” of positive charge with enough negative electrons scattered within to (COUNTERBALANCE) that positive charge.
The Structure of the Atom

William Thomson (Plum Pudding Model)

• Reasoned that the (ATOM) might be thought of as a uniform “pudding” of positive charge with enough negative electrons scattered within to (COUNTERBALANCE) that positive charge.
18. The Structure of the Atom

Ernest Rutherford (1911)

• Explained the _______ atom.
• Atom has a dense ______ of positive charge called the nucleus.
• Electrons travel around the nucleus at a relatively _____ distance.
• A ______ has the same magnitude of charge as the electron, but its charge is positive.
The Structure of the Atom

Ernest Rutherford (1911)

• Explained the (NUCLEAR) atom.
• Atom has a dense (CENTER) of positive charge called the nucleus.
• Electrons travel around the nucleus at a relatively (LARGE) distance.
• A (PROTON) has the same magnitude of charge as the electron, but its charge is positive.
19. The Structure of the Atom

Rutherford and Chadwick (1932)

• Most nuclei also contain a _______ particle called the neutron.
• A neutron is slightly more massive than a ______ but has no charge.
The Structure of the Atom

Rutherford and Chadwick (1932)

• Most nuclei also contain a (NEUTRAL) particle called the neutron.
• A neutron is slightly more massive than a (PROTON) but has no charge.
20.	Introduction to the Modern Concept of Atomic Structure

The atom contains:

•	_________ – found outside the nucleus; negatively charged
•	_______ – found in the nucleus; positive charge equal in magnitude to the electron’s negative charge
•	____
20. Introduction to the Modern Concept of Atomic Structure

The atom contains:

• _________ – found outside the nucleus; negatively charged
• _______ – found in the nucleus; positive charge equal in magnitude to the electron’s negative charge
• ________ – found in the nucleus; no charge; virtually same mass as a proton
Introduction to the Modern Concept of Atomic Structure

The atom contains:

•	(ELECTRONS) – found outside the nucleus; negatively charged
•	(PROTONS) – found in the nucleus; positive charge equal in magnitude to the electron’s negative charge
•	(NEU
Introduction to the Modern Concept of Atomic Structure

The atom contains:

• (ELECTRONS) – found outside the nucleus; negatively charged
• (PROTONS) – found in the nucleus; positive charge equal in magnitude to the electron’s negative charge
• (NEUTRONS) – found in the nucleus; no charge; virtually same mass as a proton
21. Introduction to the Modern Concept of Atomic Structure

• The nucleus is:
 _____ compared with the overall size of the atom.
 Extremely _____; accounts for almost all of the atom’s mass.
Introduction to the Modern Concept of Atomic Structure

• The nucleus is:
 (SMALL) compared with the overall size of the atom.
 Extremely (DENSE); accounts for almost all of the atom’s mass.
22.	Introduction to the Modern Concept of ______ _________
22. Introduction to the Modern Concept of ______ _________
Introduction to the Modern Concept of (ATOMIC STRUCTURE)
Introduction to the Modern Concept of (ATOMIC STRUCTURE)
23. Introduction to the Modern Concept of Atomic Structure

Why do different atoms have different chemical properties?

• The chemistry of an atom arises from its _________.
• Electrons are the parts of atoms that “intermingle” when atoms combine to form _________.
• It is the number of electrons that really determines ________ ________.
Introduction to the Modern Concept of Atomic Structure

Why do different atoms have different chemical properties?

• The chemistry of an atom arises from its (ELECTRONS).
• Electrons are the parts of atoms that “intermingle” when atoms combine to form (MOLECULES).
• It is the number of electrons that really determines (CHEMICAL BEHAVIOR).
24. Isotopes

• Atoms with the same number of protons but different numbers of ________.
• Show almost identical ________ __________; chemistry of atom is due to its electrons.
• In nature most elements contain ________ of isotopes.
Isotopes

• Atoms with the same number of protons but different numbers of (NEUTRONS).
• Show almost identical (CHEMICAL PROPERTIES); chemistry of atom is due to its electrons.
• In nature most elements contain (MIXTURES) of isotopes.
25.	Isotopes

Two ________ of Sodium
25. ________

Two ________ of Sodium
Isotopes

Two (ISOTOPES) of Sodium
(ISOTOPES)

Two (ISOTOPES) of Sodium
1.	Isotopes

•	X = the symbol of the _______
•	Z = the atomic number (# of _______)
•	A = the mass number (# of _______ and ________)
1. Isotopes

• X = the symbol of the _______
• Z = the atomic number (# of _______)
• A = the mass number (# of _______ and ________)
Isotopes

•	X = the symbol of the (ELEMENT)
•	Z = the atomic number (# of [PROTONS])
•	A = the mass number (# of [PROTONS] and [NEUTRONS])
Isotopes

• X = the symbol of the (ELEMENT)
• Z = the atomic number (# of [PROTONS])
• A = the mass number (# of [PROTONS] and [NEUTRONS])
1.	Isotopes

Isotopes – An Example

•	C = the symbol for ______
•	6 = the ______ ______ (6 protons)
•	14 = the ____ ______ (6 protons and 8 neutrons)
1. Isotopes

Isotopes – An Example

• C = the symbol for ______
• 6 = the ______ ______ (6 protons)
• 14 = the ____ ______ (6 protons and 8 neutrons)
Isotopes – An Example

•	C = the symbol for (CARBON)
•	6 = the (ATOMIC NUMBER) (6 protons)
•	14 = the (MASS NUMBER) (6 protons and 8 neutrons)
Isotopes – An Example

• C = the symbol for (CARBON)
• 6 = the (ATOMIC NUMBER) (6 protons)
• 14 = the (MASS NUMBER) (6 protons and 8 neutrons)
1.	Isotopes

Isotopes – An Example

•	C = the symbol for carbon
•	6 = the atomic number (_ _______)
•	12 = the mass number (_ _______ and _ ________)
1. Isotopes

Isotopes – An Example

• C = the symbol for carbon
• 6 = the atomic number (_ _______)
• 12 = the mass number (_ _______ and _ ________)
Isotopes – An Example

•	C = the symbol for carbon
•	6 = the atomic number ([6 PROTONS])
•	12 = the mass number ([6 PROTONS AND 6 NEUTRONS])
Isotopes – An Example

• C = the symbol for carbon
• 6 = the atomic number ([6 PROTONS])
• 12 = the mass number ([6 PROTONS AND 6 NEUTRONS])
29. Isotopes

Exercise

A certain isotope X contains 23 protons and 28 neutrons.
• What is the mass number of this isotope?
• Identify the element.
Isotopes

Exercise

A certain isotope X contains 23 protons and 28 neutrons.
• What is the mass number of this isotope?
Identify the element.

Mass Number = 51
Vanadium
30.	Introduction to the Periodic Table

The Periodic Table

•	The periodic table shows all of the known elements in order of increasing ______ ______.
30. Introduction to the Periodic Table

The Periodic Table

• The periodic table shows all of the known elements in order of increasing ______ ______.
The Periodic Table

•	The periodic table shows all of the known elements in order of increasing (ATOMIC NUMBER).
The Periodic Table

• The periodic table shows all of the known elements in order of increasing (ATOMIC NUMBER).
31. Introduction to the Periodic Table

The Periodic Table

• Metals vs. Nonmetals
• Groups or Families – elements in the same vertical columns; have similar ________ __________
• _______ – horizontal rows of elements
The Periodic Table

• Metals vs. Nonmetals
• Groups or Families – elements in the same vertical columns; have similar (CHEMICAL PROPERTIES)
• (PERIODS) – horizontal rows of elements
32.	Introduction to the Periodic Table

The Periodic Table

•	Most elements are ______ and occur on the left side.
•	The _________ appear on the right side.
•	__________ are elements that have some metallic and some nonmetallic properties.
32. Introduction to the Periodic Table

The Periodic Table

• Most elements are ______ and occur on the left side.
• The _________ appear on the right side.
• __________ are elements that have some metallic and some nonmetallic properties.
The Periodic Table

•	Most elements are (METALS) and occur on the left side.
•	The (NONMETALS) appear on the right side.
•	(METALLOIDS) are elements that have some metallic and some nonmetallic properties.
The Periodic Table

• Most elements are (METALS) and occur on the left side.
• The (NONMETALS) appear on the right side.
• (METALLOIDS) are elements that have some metallic and some nonmetallic properties.
33. Introduction to the Periodic Table

Physical Properties of Metals

1. Efficient __________ of heat and electricity
2. ____________ (they can be hammered into thin sheets)
3. _________ (they can be pulled into wires)
4. A lustrous (shiny) appearance
Physical Properties of Metals

1. Efficient (CONDUCTION) of heat and electricity
2. (MALLEABILITY) (they can be hammered into thin sheets)
3. (DUCTILITY) (they can be pulled into wires)
4. A lustrous (shiny) appearance
1. Natural States of the Elements

• Most elements are very ________.
• Elements are not generally found in __________ form.
 Exceptions are:
• Noble metals – gold, platinum and silver
• Noble gases – Group 8
Natural States of the Elements

• Most elements are very (REACTIVE).
• Elements are not generally found in (UNCOMBINED) form.
 Exceptions are:
• Noble metals – gold, platinum and silver
• Noble gases – Group 8
1.	Natural States of the Elements

Diatomic Molecules

________ gas contains N2 molecules.
1. Natural States of the Elements

Diatomic Molecules

________ gas contains N₂ molecules.
Natural States of the Elements

Diatomic Molecules

(NITROGEN) gas contains N2 molecules.
Natural States of the Elements

Diatomic Molecules

(NITROGEN) gas contains N₂ molecules.
1.	Natural States of the Elements

Diatomic Molecules

______ gas contains O2 molecules.
1. Natural States of the Elements

Diatomic Molecules

______ gas contains O₂ molecules.
Natural States of the Elements

Diatomic Molecules

(OXYGEN) gas contains O2 molecules.
Natural States of the Elements

Diatomic Molecules

(OXYGEN) gas contains O₂ molecules.
4.	Natural States of the Elements

Diatomic Molecules
4. Natural States of the Elements

________ _________
Natural States of the Elements

Diatomic Molecules
Natural States of the Elements

(DIATOMIC MOLECULES)
5. Natural States of the Elements

Allotropes

• Different _____ of a given element.
• Example:
 Solid carbon occurs in three _____.
• Diamond
• Graphite
• Buckminsterfullerene
Natural States of the Elements

Allotropes

• Different (FORMS) of a given element.
• Example:
 Solid carbon occurs in three (FORMS).
• Diamond
• Graphite
• Buckminsterfullerene
6.	Natural States of the Elements

Carbon __________
6. Natural States of the Elements

Carbon __________
Natural States of the Elements

Carbon (ALLOTROPES)
Natural States of the Elements

Carbon (ALLOTROPES)
7.	Ions

•	_____ can form ions by gaining or losing electrons.
	Metals tend to lose one or more electrons to form positive ions called _______.
	_______ are generally named by using the name of the parent atom.
7. Ions

• _____ can form ions by gaining or losing electrons.
 Metals tend to lose one or more electrons to form positive ions called _______.
 _______ are generally named by using the name of the parent atom.
Ions

•	(ATOMS) can form ions by gaining or losing electrons.
	Metals tend to lose one or more electrons to form positive ions called (CATIONS).
	(CATIONS) are generally named by using the name of the parent atom.
Ions

• (ATOMS) can form ions by gaining or losing electrons.
 Metals tend to lose one or more electrons to form positive ions called (CATIONS).
 (CATIONS) are generally named by using the name of the parent atom.
8.	Ions

•	_________ tend to gain one or more electrons to form negative ions called ______.
•	Anions are named by using the root of the atom name followed by the suffix –ide.
8. Ions

• _________ tend to gain one or more electrons to form negative ions called ______.
• ______ are named by using the root of the atom name followed by the suffix –ide.
Ions

•	(NONMETALS) tend to gain one or more electrons to form negative ions called (ANIONS).
•	Anions are named by using the root of the atom name followed by the suffix –ide.
Ions

• (NONMETALS) tend to gain one or more electrons to form negative ions called (ANIONS).
• (ANIONS) are named by using the root of the atom name followed by the suffix –ide.
9.	Ions

Ion Charges and the Periodic Table

•	The ___ that a particular atom will form can be predicted from the periodic table.
9. ____

___ Charges and the Periodic Table

• The ___ that a particular atom will form can be predicted from the periodic table.
Ions

Ion Charges and the Periodic Table

•	The (ION) that a particular atom will form can be predicted from the periodic table.
(IONS)

(ION) Charges and the Periodic Table

• The (ION) that a particular atom will form can be predicted from the periodic table.
10.	Ions

Ion _______ and the Periodic Table
10. Ions

Ion _______ and the Periodic Table
Ions

Ion (CHARGES) and the Periodic Table
Ions

Ion (CHARGES) and the Periodic Table
11. Ions

Exercise

An ion with a 3+ charge contains 23 electrons. Which ion is it?

a) Fe³⁺
b) V³⁺
c) Ca³⁺
d) Sc³⁺
Ions

Exercise

An ion with a 3+ charge contains 23 electrons. Which ion is it?

a) Fe³⁺
12. Ions

Exercise

A certain ion X⁺ contains 54 electrons and 78 neutrons. What is the mass number of this ion?
Exercise

A certain ion X⁺ contains 54 electrons and 78 neutrons. What is the mass number of this ion?

133
13.	Compounds That Contain Ions

•	____ combine to form ionic compounds. 
•	Properties of ionic compounds
	High _______ ______
	_______ electricity 
•	If melted
•	If dissolved in water
13. Compounds That Contain Ions

• ____ combine to form ionic compounds.
• Properties of ionic compounds
 High _______ ______
 _______ electricity
• If melted
• If dissolved in water
Compounds That Contain Ions

•	(IONS) combine to form ionic compounds. 
•	Properties of ionic compounds
	High (MELTING POINTS)
	(CONDUCT) electricity 
•	If melted
•	If dissolved in water
Compounds That Contain Ions

• (IONS) combine to form ionic compounds.
• Properties of ionic compounds
 High (MELTING POINTS)
 (CONDUCT) electricity
• If melted
• If dissolved in water
14. Compounds That Contain Ions

• Ionic compounds are electrically _______.
• The _______ on the anions and cations in the compound must sum to zero.
Compounds That Contain Ions

• Ionic compounds are electrically (NEUTRAL).
• The (CHARGES) on the anions and cations in the compound must sum to zero.
15.	Compounds That Contain Ions

Formulas for Ionic Compounds

•	Write the cation element symbol followed by the anion element symbol.
•	The number of cations and anions must be correct for their charges to sum to ____.
15. Compounds That Contain Ions

Formulas for Ionic Compounds

• Write the cation element symbol followed by the anion element symbol.
• The number of cations and anions must be correct for their charges to sum to ____.
Compounds That Contain Ions

Formulas for Ionic Compounds

•	Write the cation element symbol followed by the anion element symbol.
•	The number of cations and anions must be correct for their charges to sum to (ZERO).
Compounds That Contain Ions

Formulas for Ionic Compounds

• Write the cation element symbol followed by the anion element symbol.
• The number of cations and anions must be correct for their charges to sum to (ZERO).
16. Compounds That Contain Ions

Concept Check

A compound contains an unknown ion X and has the formula XCl₂. Ion X contains 20 electrons. What is the identity of X?

a) Ti²⁺
b) Sc⁺
c) Ca²⁺
d) Cr²⁺
Compounds That Contain Ions

Concept Check

A compound contains an unknown ion X and has the formula XCl₂. Ion X contains 20 electrons. What is the identity of X?

a) Ti²⁺
17. Compounds That Contain Ions

Concept Check

A member of the alkaline earth metal family whose most stable ion contains 36 electrons forms a compound with bromine. What is the correct formula for this compound?

a) CaBr₂
b) KrBr
c) RbBr
d) SrBr₂
Compounds That Contain Ions

Concept Check

A member of the alkaline earth metal family whose most stable ion contains 36 electrons forms a compound with bromine. What is the correct formula for this compound?

d) SrBr₂
A two-element compound.
1. Binary Compound
A two-element compound consisting of a cation and an anion.
2. Binary Ionic Compound
An ion containing a number of atoms.
3. Polyatomic Ion
A polyatomic ion containing at least one oxygen atom and one or more atoms of at least one other element.
4. Oxyanion
A substance that produces hydrogen ions in aqueous solution; proton donor.
5. Acid
Radiant energy that exhibits wavelike behavior and travels through space at the speed of light in a vacuum.
1. Electromagnetic Radiation
The distance between two consecutive peaks or troughs in a wave.
2. Wavelength
The number of waves (cycles) per second that pass a given point in space.
3. Frequency
A “particle” of electromagnetic radiation.
4. Photon
Energy levels where only certain values are allowed.
5. Quantized Energy Levels
The waves produced by an electron confined in its orbit about the nucleus sets up a "standing wave",( a specific number of "bounces" each second), of specific wavelength, energy and frequency (i.e., Bohr's energy levels) like a rubber band when stretched and released (online definition).
6. Wave Mechanical Model
A representation of the space occupied by an electron in an atom; the probability distribution for the electron.
7. Orbital
Discrete energy levels.
8. Principal Energy Levels
Subdivision of the principal energy level.
9. Sublevels
In a given atom, no two elements can occupy the same atomic orbital and have the same spin.
10. Pauli Exclusion Principle
The distribution of electrons of an atom or molecule (or other physical structure) in atomic or molecular orbitals (online definition).
11. Electron Configuration
Consists of a box representing each orbital and a half arrow representing each electron; can be read directly from the periodic table as well as electron configurations (online definition).
12. Orbital (Box) Diagram
The electrons in the outermost occupied principal quantum level of an atom.
13. Valence Electrons
An inner electron in an atom; one that is not in the outermost (valence) principal quantum level.
14. Core Electrons
A group of fourteen elements following lanthanum on the periodic table, in which the 41 orbitals are being filled.
15. Lanthanide Series
A group of fourteen elements following actinium on the periodic table, in which the 51 orbitals are being filled.
16. Actinide Series
Elements in the groups labeled 1, 2, 3, 4, 5, 6, 7, and 8 on the periodic table. The group number gives the sum of the valence s and p electrons.
17. Main-Group (Representative) Elements
An element that gives up electrons relatively easily and is typically lustrous, malleable, and a good conductor of heat and electricity.
18. Metals
An element that does not exhibit metallic characteristics; chemically, a typical ________ accepts electrons from a metal.
19. Nonmetals
An element that has both metallic and nonmetallic properties.
20. Metalloids
The quantity of energy required to remove an electron from a gaseous atom or ion.
21. Ionization Energy
Each atom has a nucleus inside and electrons zooming around outside the nucleus. The size of an atom depends on how far away its outermost (valence) electrons are from the nucleus. If they are very close to the nucleus, the atom will be very small. If they are far away, the atom will be quite a bit larger. So the ______ ____ is determined by how much space the electrons take up (online definition).
22. Atomic Size
1. Rutherford’s Atom

Nuclear Model of the Atom

• The atom has a small dense nucleus which
 is positively charged.
 contains _______ (+1 charge).
 contains ________ (no charge).
• The remainder of the atom
 is mostly empty space.
 contains _________ (–1 charge).
Rutherford’s Atom

Nuclear Model of the Atom

• The atom has a small dense nucleus which
 is positively charged.
 contains (PROTONS) (+1 charge).
 contains (NEUTRONS) (no charge).
• The remainder of the atom
 is mostly empty space.
 contains (ELECTRONS) (–1 charge).
2.	Rutherford’s Atom

•	The nuclear charge (n+) is balanced by the presence of n electrons moving in some way around the _______.
•	What are the electrons doing?
•	How are the electrons arranged and how do they move?
2. Rutherford’s Atom

• The nuclear charge (n+) is balanced by the presence of n electrons moving in some way around the _______.
• What are the electrons doing?
• How are the electrons arranged and how do they move?
Rutherford’s Atom

•	The nuclear charge (n+) is balanced by the presence of n electrons moving in some way around the _______.
•	What are the electrons doing?
•	How are the electrons arranged and how do they move?
Rutherford’s Atom

• The nuclear charge (n+) is balanced by the presence of n electrons moving in some way around the (NUCLEUS).
• What are the electrons doing?
• How are the electrons arranged and how do they move?
3.	Electromagnetic Radiation

One of the ways that ______ travels through
3. Electromagnetic Radiation

One of the ways that ______ travels through
Electromagnetic Radiation

One of the ways that (ENERGY) travels through
Electromagnetic Radiation

One of the ways that (ENERGY) travels through
4.	Electromagnetic Radiation

Characteristics

•	Wavelength (λ) – distance between two _____ or _______ in a wave.
4. Electromagnetic Radiation

Characteristics

• Wavelength (λ) – distance between two _____ or _______ in a wave.
Electromagnetic Radiation

Characteristics

•	Wavelength (λ) – distance between two (PEAKS) or (TROUGHS) in a wave.
Electromagnetic Radiation

Characteristics

• Wavelength (λ) – distance between two (PEAKS) or (TROUGHS) in a wave.
5.	Electromagnetic Radiation

Characteristics

•	Frequency (v) – number of _____ (cycles) per second that pass a given point in space
•	Speed (c) – speed of light (2.9979×108 m/s)
5. Electromagnetic Radiation

Characteristics

• Frequency (v) – number of _____ (cycles) per second that pass a given point in space
• Speed (c) – speed of light (2.9979×108 m/s)
Electromagnetic Radiation

Characteristics

•	Frequency (v) – number of (WAVES) (cycles) per second that pass a given point in space
•	Speed (c) – speed of light (2.9979×108 m/s)
Electromagnetic Radiation

Characteristics

• Frequency (v) – number of (WAVES) (cycles) per second that pass a given point in space
• Speed (c) – speed of light (2.9979×108 m/s)
6.	Electromagnetic Radiation

Dual Nature of Light

•	Wave
•	______ – packet of energy
6. Electromagnetic Radiation

Dual Nature of Light

• Wave
• ______ – packet of energy
Electromagnetic Radiation

Dual Nature of Light

•	Wave
•	(PHOTON) – packet of energy
Electromagnetic Radiation

Dual Nature of Light

• Wave
• (PHOTON) – packet of energy
7.	Electromagnetic Radiation

Different Wavelengths Carry Different Amounts of ______
7. Electromagnetic Radiation

Different Wavelengths Carry Different Amounts of ______
Electromagnetic Radiation

Different Wavelengths Carry Different Amounts of (ENERGY)
Electromagnetic Radiation

Different Wavelengths Carry Different Amounts of (ENERGY)
8.	Emission of Energy by Atoms

•	Atoms can give off _____. 
	They first must _______ energy and become excited. 
	The energy is released in the form of a photon. 
	The energy of the photon corresponds exactly to the energy change experienced by t
8. Emission of Energy by Atoms

• Atoms can give off _____.
 They first must _______ energy and become excited.
 The energy is released in the form of a photon.
 The energy of the photon corresponds exactly to the energy change experienced by the emitting ____.
Emission of Energy by Atoms

•	Atoms can give off (LIGHT). 
	They first must (RECEIVE) energy and become excited. 
	The energy is released in the form of a photon. 
	The energy of the photon corresponds exactly to the energy change experienced by
Emission of Energy by Atoms

• Atoms can give off (LIGHT).
 They first must (RECEIVE) energy and become excited.
 The energy is released in the form of a photon.
 The energy of the photon corresponds exactly to the energy change experienced by the emitting (ATOM).
9.	The Energy Levels of Hydrogen

•	Atomic states 
	Excited state – atom with ______ energy 
	Ground state – atom in the ______ ________ state 
•	When an H atom absorbs energy from an outside source it enters an _______ state.
9. The Energy Levels of Hydrogen

• Atomic states
 Excited state – atom with ______ energy
 Ground state – atom in the ______ ________ state
• When an H atom absorbs energy from an outside source it enters an _______ state.
The Energy Levels of Hydrogen

•	Atomic states 
	Excited state – atom with (EXCESS) energy 
	Ground state – atom in the (LOWEST POSSIBLE) state 
•	When an H atom absorbs energy from an outside source it enters an (EXCITED) state.
The Energy Levels of Hydrogen

• Atomic states
 Excited state – atom with (EXCESS) energy
 Ground state – atom in the (LOWEST POSSIBLE) state
• When an H atom absorbs energy from an outside source it enters an (EXCITED) state.
10.	The Energy Levels of Hydrogen

Energy Level Diagram

•	Energy in the photon corresponds to the energy used by the ____ to get to the excited state.
10. The Energy Levels of Hydrogen

Energy Level Diagram

• Energy in the photon corresponds to the energy ____ by the ____ to get to the excited state.
The Energy Levels of Hydrogen

Energy Level Diagram

•	Energy in the photon corresponds to the energy used by the (ATOM) to get to the excited state.
The Energy Levels of Hydrogen

Energy Level Diagram

• Energy in the photon corresponds to the energy (USED) by the (ATOM) to get to the excited state.
11.	The Energy Levels of Hydrogen

•	Only certain types of _______ are produced when H atoms release energy.  Why?
11. The Energy Levels of Hydrogen

• Only certain types of _______ are produced when H atoms release energy. Why?
The Energy Levels of Hydrogen

•	Only certain types of (PHOTONS) are produced when H atoms release energy.  Why?
The Energy Levels of Hydrogen

• Only certain types of (PHOTONS) are produced when H atoms release energy. Why?
12.	The Energy Levels of Hydrogen

Quantized Energy Levels

•	Since only certain energy changes occur the H atom must contain ________ energy levels.
12. The Energy Levels of Hydrogen

Quantized Energy Levels

• Since only certain energy changes occur the H atom must contain ________ energy levels.
The Energy Levels of Hydrogen

Quantized Energy Levels

•	Since only certain energy changes occur the H atom must contain (DISCRETE) energy levels.
The Energy Levels of Hydrogen

Quantized Energy Levels

• Since only certain energy changes occur the H atom must contain (DISCRETE) energy levels.
13.	The Energy Levels of Hydrogen

Quantized Energy Levels

•	The energy levels of ___ atoms are quantized.
13. The Energy Levels of Hydrogen

Quantized Energy Levels

• The energy levels of ___ atoms are quantized.
The Energy Levels of Hydrogen

Quantized Energy Levels

•	The energy levels of (ALL) atoms are quantized.
The Energy Levels of Hydrogen

Quantized Energy Levels

• The energy levels of (ALL) atoms are quantized.
14. The Energy Levels of Hydrogen

Concept Check

Why is it significant that the color emitted from the hydrogen emission spectrum is not white?

How does the emission spectrum support the idea of quantized energy levels?
The Energy Levels of Hydrogen

Concept Check

Why is it significant that the color emitted from the hydrogen emission spectrum is not white?

How does the emission spectrum support the idea of quantized energy levels?
15. The Energy Levels of Hydrogen

Concept Check

When an electron is excited in an atom or ion

a) only specific quantities of energy are released in order for the electron to return to its ground state.
b) white light is never observed when the electron returns to its ground state.
c) the electron is only excited to certain energy levels.
d) All of the above statements are true when an electron is excited.
The Energy Levels of Hydrogen

Concept Check

When an electron is excited in an atom or ion

d) All of the above statements are true when an electron is excited.
16.	The Bohr Model of the Atom

•	Quantized energy levels
•	Electron moves in a circular orbit.
•	Electron jumps between levels by absorbing or emitting a photon of a particular wavelength.
16. The Bohr Model of the Atom

• _________ energy levels
• Electron moves in a circular orbit.
• Electron jumps between levels by _________ or ________ a photon of a particular wavelength.
The Bohr Model of the Atom

•	Quantized energy levels
•	Electron moves in a circular orbit.
•	Electron jumps between levels by absorbing or emitting a photon of a particular wavelength.
The Bohr Model of the Atom

• (QUANTIZED) energy levels
• Electron moves in a circular orbit.
• Electron jumps between levels by (ABSORBING) or (EMITTING) a photon of a particular wavelength.
17. The Bohr Model of the Atom

• Bohr’s model of the atom was _________.
 Electron does not move in a ________ orbit.
The Bohr Model of the Atom

• Bohr’s model of the atom was (INCORRECT).
 Electron does not move in a (CIRCULAR) orbit.
18.	The Wave Mechanical Model of the Atom

Orbitals

•	Nothing like ______.
•	Probability of finding the electron within a certain _____.
•	This model gives no information about when the electron occupies a certain point in space or how it _____.
18. The Wave Mechanical Model of the Atom

Orbitals

• Nothing like ______.
• Probability of finding the electron within a certain _____.
• This model gives no information about when the electron occupies a certain point in space or how it _____.
The Wave Mechanical Model of the Atom

Orbitals

•	Nothing like (ORBITS).
•	Probability of finding the electron within a certain (SPACE).
•	This model gives no information about when the electron occupies a certain point in space or how it (MOVES).
The Wave Mechanical Model of the Atom

Orbitals

• Nothing like (ORBITS).
• Probability of finding the electron within a certain (SPACE).
• This model gives no information about when the electron occupies a certain point in space or how it (MOVES).
19.	The Hydrogen Orbitals

Orbitals

•	Orbitals do not have sharp __________.
•	Chemists arbitrarily define an orbital’s size as the sphere that contains __% of the total electron probability.
19. The Hydrogen Orbitals

Orbitals

• Orbitals do not have sharp __________.
• Chemists arbitrarily define an orbital’s size as the sphere that contains __% of the total electron probability.
The Hydrogen Orbitals

Orbitals

•	Orbitals do not have sharp (BOUNDARIES).
•	Chemists arbitrarily define an orbital’s size as the sphere that contains (90)% of the total electron probability.
The Hydrogen Orbitals

Orbitals

• Orbitals do not have sharp (BOUNDARIES).
• Chemists arbitrarily define an orbital’s size as the sphere that contains (90)% of the total electron probability.
20.	The Hydrogen Orbitals

Hydrogen Energy Levels

•	Hydrogen has ________ energy levels which are called principal energy levels and are labeled with _____ numbers.
20. The Hydrogen Orbitals

Hydrogen Energy Levels

• Hydrogen has ________ energy levels which are called principal energy levels and are labeled with _____ numbers.
The Hydrogen Orbitals

Hydrogen Energy Levels

•	Hydrogen has (DISCRETE) energy levels which are called principal energy levels and are labeled with (WHOLE) numbers
The Hydrogen Orbitals

Hydrogen Energy Levels

• Hydrogen has (DISCRETE) energy levels which are called principal energy levels and are labeled with (WHOLE) numbers
21.	The Hydrogen Orbitals

Hydrogen Energy Levels

•	Each principal energy level is divided into _________ which are labeled with numbers and letters which indicate the _____ of the orbital.
21. The Hydrogen Orbitals

Hydrogen Energy Levels

• Each principal energy level is divided into _________ which are labeled with numbers and letters which indicate the _____ of the orbital.
The Hydrogen Orbitals

Hydrogen Energy Levels

•	Each principal energy level is divided into (SUBLEVELS) which are labeled with numbers and letters which indicate the (SHAPE) of the orbital.
The Hydrogen Orbitals

Hydrogen Energy Levels

• Each principal energy level is divided into (SUBLEVELS) which are labeled with numbers and letters which indicate the (SHAPE) of the orbital.
22.	The Hydrogen Orbitals

Hydrogen Energy Levels

•	The s and p types of ________.
22. The Hydrogen Orbitals

Hydrogen Energy Levels

• The s and p types of ________.
The Hydrogen Orbitals

Hydrogen Energy Levels

•	The s and p types of (SUBLEVEL).
The Hydrogen Orbitals

Hydrogen Energy Levels

• The s and p types of (SUBLEVEL).
23. The Hydrogen Orbitals

Orbital Labels

1. The number tells the principal energy level.
2. The letter tells the shape.
 The letter s means a _________ orbital.
 The letter p means a ___-_____ orbital. The x, y, or z subscript on a p orbital label tells along which of the coordinate ____ the two lobes lie.
The Hydrogen Orbitals

Orbital Labels

1. The number tells the principal energy level.
2. The letter tells the shape.
 The letter s means a (SPHERICAL) orbital.
 The letter p means a (TWO–LOBED) orbital. The x, y, or z subscript on a p orbital label tells along which of the coordinate (AXES) the two lobes lie.
24. The Hydrogen Orbitals

Hydrogen Orbitals

• Why does an H atom have so many orbitals and only 1 electron?
 An orbital is a potential _____ for an electron.
 Atoms can have ____ potential orbitals.
The Hydrogen Orbitals

Hydrogen Orbitals

• Why does an H atom have so many orbitals and only 1 electron?
 An orbital is a potential (SPACE) for an electron.
 Atoms can have (MANY) potential orbitals.
25. The Wave Mechanical Model: Further Development

Atoms Beyond Hydrogen

• The Bohr model was discarded because it does not apply to ___ atoms.
• Atoms beyond hydrogen have an _____ number of protons and electrons which need one more property to determine how the electrons are arranged.
 Spin – ________ spins like a top.
The Wave Mechanical Model: Further Development

Atoms Beyond Hydrogen

• The Bohr model was discarded because it does not apply to (ALL) atoms.
• Atoms beyond hydrogen have an (EQUAL) number of protons and electrons which need one more property to determine how the electrons are arranged.
 Spin – (ELECTRON) spins like a top.
26. The Wave Mechanical Model: Further Development

Atoms Beyond Hydrogen

• Pauli Exclusion Principle – an atomic orbital can hold a _______ of 2 electrons and those 2 electrons must have ________ spins.
The Wave Mechanical Model: Further Development

Atoms Beyond Hydrogen

• Pauli Exclusion Principle – an atomic orbital can hold a (MAXIMUM) of 2 electrons and those 2 electrons must have (OPPOSITE) spins.
27. The Wave Mechanical Model: Further Development

Principal Components of the Wave Mechanical Model of the Atom

1. _____ have a series of energy levels called principal energy levels (n = 1, 2, 3, etc.).
2. The energy of the level increases as the value of n increases.
3. Each principal energy level contains one or more types of ________, called sublevels.
4. The number of sublevels present in a given principal energy level ______ n.
The Wave Mechanical Model: Further Development

Principal Components of the Wave Mechanical Model of the Atom

1. (ATOMS) have a series of energy levels called principal energy levels (n = 1, 2, 3, etc.).
2. The energy of the level increases as the value of n increases.
3. Each principal energy level contains one or more types of (ORBITALS), called sublevels.
4. The number of sublevels present in a given principal energy level (EQUALS) n.
28. The Wave Mechanical Model: Further Development

Principal Components of the Wave Mechanical Model of the Atom

5. The n value is always used to label the orbitals of a given principal level and is followed by a ______ that indicates the type (shape) of the orbital (1s, 3p, etc.).
6. An orbital can be _____ or it can contain one or two electrons, but never more than two. If two electrons occupy the same orbital, they must have ________ spins.
The Wave Mechanical Model: Further Development

Principal Components of the Wave Mechanical Model of the Atom

5. The n value is always used to label the orbitals of a given principal level and is followed by a (LETTER) that indicates the type (shape) of the orbital (1s, 3p, etc.).
6. An orbital can be (EMPTY) or it can contain one or two electrons, but never more than two. If two electrons occupy the same orbital, they must have (OPPOSITE) spins.
29. The Wave Mechanical Model: Further Development

Principal Components of the Wave Mechanical Model of the Atom

7. The shape of an orbital does not indicate the details of electron ________. It indicates the ___________ distribution for an electron residing in that orbital.
The Wave Mechanical Model: Further Development

Principal Components of the Wave Mechanical Model of the Atom

7. The shape of an orbital does not indicate the details of electron (MOVEMENT). It indicates the (PROBABILITY) distribution for an electron residing in that orbital.
30. The Wave Mechanical Model: Further Development

Concept Check

Which of the following statements best describes the movement of electrons in a p orbital?

a) The electron movement cannot be exactly determined.
b) The electrons move within the two lobes of the p orbital, but never beyond the outside surface of the orbital.
c) The electrons are concentrated at the center (node) of the two lobes.
d) The electrons move along the outer surface of the p orbital, similar to a “figure 8” type of movement.
The Wave Mechanical Model: Further Development

Concept Check

Which of the following statements best describes the movement of electrons in a p orbital?

a) The electron movement cannot be exactly determined.
31.	Electron Arrangements in the First Eighteen Atoms on the Periodic Table

H Atom

•	Electron configuration – electron arrangement 	1s1
•	Orbital diagram – orbital is a box grouped by ________ containing _____(s) to represent electrons
31. Electron Arrangements in the First Eighteen Atoms on the Periodic Table

H Atom

• Electron configuration – electron arrangement 1s¹
• Orbital diagram – orbital is a box grouped by ________ containing _____(s) to represent electrons
Electron Arrangements in the First Eighteen Atoms on the Periodic Table

H Atom

•	Electron configuration – electron arrangement 	1s1
•	Orbital diagram – orbital is a box grouped by (SUBLEVEL) containing (ARROW)(s) to represent electrons
Electron Arrangements in the First Eighteen Atoms on the Periodic Table

H Atom

• Electron configuration – electron arrangement 1s¹
• Orbital diagram – orbital is a box grouped by (SUBLEVEL) containing (ARROW)(s) to represent electrons
32.	Electron Arrangements in the First Eighteen Atoms on the Periodic Table

Li Atom

•	Electron _____________	
		1s2 2s1
•	Orbital diagram
32. Electron Arrangements in the First Eighteen Atoms on the Periodic Table

Li Atom

• Electron _____________
1s² 2s¹
• Orbital diagram
Electron Arrangements in the First Eighteen Atoms on the Periodic Table

Li Atom

•	Electron (CONFIGURATION)	
		1s2 2s1
•	Orbital diagram
Electron Arrangements in the First Eighteen Atoms on the Periodic Table

Li Atom

• Electron (CONFIGURATION)
1s² 2s¹
• Orbital diagram
33.	Electron Arrangements in the First Eighteen Atoms on the Periodic Table

O Atom

•	The lowest energy configuration for an atom is the one having the maximum number of ________ electrons in a particular set of degenerate (same energy) orbitals.
33. Electron Arrangements in the First Eighteen Atoms on the Periodic Table

O Atom

• The lowest energy configuration for an atom is the one having the maximum number of ________ electrons in a particular set of degenerate (same energy) orbitals.
Electron Arrangements in the First Eighteen Atoms on the Periodic Table

O Atom

•	The lowest energy configuration for an atom is the one having the maximum number of (UNPAIRED) electrons in a particular set of degenerate (same energy) orbitals.
Electron Arrangements in the First Eighteen Atoms on the Periodic Table

O Atom

• The lowest energy configuration for an atom is the one having the maximum number of (UNPAIRED) electrons in a particular set of degenerate (same energy) orbitals.
34.	Electron Arrangements in the First Eighteen Atoms on the Periodic Table

•	The electron configurations in the sublevel last ________ for the first eighteen elements.
34. Electron Arrangements in the First Eighteen Atoms on the Periodic Table

• The electron configurations in the sublevel last ________ for the first eighteen elements.
Electron Arrangements in the First Eighteen Atoms on the Periodic Table

•	The electron configurations in the sublevel last (OCCUPIED) for the first eighteen elements.
Electron Arrangements in the First Eighteen Atoms on the Periodic Table

• The electron configurations in the sublevel last (OCCUPIED) for the first eighteen elements.
35. Electron Arrangements in the First Eighteen Atoms on the Periodic Table

Classifying Electrons

• Core electrons – _____ electrons.
• Valence electrons – electrons in the _________ (highest) principal energy level of an atom.
 1s²2s²2p⁶ (valence electrons = 8)
 The elements in the same _____ on the periodic table have the same valence electron configuration.
 Elements with the same valence electron arrangement show very similar ________ behavior.
Electron Arrangements in the First Eighteen Atoms on the Periodic Table

Classifying Electrons

• Core electrons – (INNER) electrons.
• Valence electrons – electrons in the (OUTERMOST) (highest) principal energy level of an atom.
 1s²2s²2p⁶ (valence electrons = 8)
 The elements in the same (GROUP) on the periodic table have the same valence electron configuration.
 Elements with the same valence electron arrangement show very similar (CHEMICAL) behavior.
36. Electron Arrangements in the First Eighteen Atoms on the Periodic Table

Concept Check

How many unpaired electrons does the element cobalt (Co) have in its lowest energy state?

a) 0
b) 2
c) 3
d) 7
Electron Arrangements in the First Eighteen Atoms on the Periodic Table

Concept Check

How many unpaired electrons does the element cobalt (Co) have in its lowest energy state?

c) 3
37. Electron Arrangements in the First Eighteen Atoms on the Periodic Table

Concept Check

Can an electron in a phosphorus atom ever be in a 3d orbital? Choose the best answer.

a) Yes. An electron can be excited into a 3d orbital.
b) Yes. A ground-state electron in phosphorus is located in a 3d orbital.
c) No. Only transition metal atoms can have electrons located in the d orbitals.
d) No. This would not correspond to phosphorus’ electron arrangement in its ground state.
Electron Arrangements in the First Eighteen Atoms on the Periodic Table

Concept Check

Can an electron in a phosphorus atom ever be in a 3d orbital? Choose the best answer.

a) Yes. An electron can be excited into a 3d orbital.
38.	Electron Configurations and the Periodic Table

•	Look at electron ______________ for K through Kr.
38. Electron Configurations and the Periodic Table

• Look at electron ______________ for K through Kr.
Electron Configurations and the Periodic Table

•	Look at electron (CONFIGURATIONS) for K through Kr.
Electron Configurations and the Periodic Table

• Look at electron (CONFIGURATIONS) for K through Kr.
39.	Electron Configurations and the Periodic Table

Orbital _______ and the Periodic Table
39. Electron Configurations and the Periodic Table

Orbital _______ and the Periodic Table
Electron Configurations and the Periodic Table

Orbital (FILLING) and the Periodic Table
Electron Configurations and the Periodic Table

Orbital (FILLING) and the Periodic Table
40. Electron Configurations and the Periodic Table

Orbital Filling

1. In a principal energy level that has d orbitals, the s orbital from the next level fills ______ the d orbitals in the current level.
2. After lanthanum, which has the electron configuration [Xe]6s²5d¹, a group of fourteen elements called the lanthanide series, or the lanthanides, occurs. This series of elements corresponds to the _______ of the seven 4f orbitals.
Electron Configurations and the Periodic Table

Orbital Filling

1. In a principal energy level that has d orbitals, the s orbital from the next level fills (BEFORE) the d orbitals in the current level.
2. After lanthanum, which has the electron configuration [Xe]6s²5d¹, a group of fourteen elements called the lanthanide series, or the lanthanides, occurs. This series of elements corresponds to the (FILLING) of the seven 4f orbitals.
41. Electron Configurations and the Periodic Table

Orbital Filling

3. After actinum, which has the configuration [Rn]7s²6d¹, a group of fourteen elements called the actinide series, or actinides, occurs. This series corresponds to the filling of the seven 5f ________.
Electron Configurations and the Periodic Table

Orbital Filling

3. After actinum, which has the configuration [Rn]7s²6d¹, a group of fourteen elements called the actinide series, or actinides, occurs. This series corresponds to the filling of the seven 5f (ORBITALS).
42. Electron Configurations and the Periodic Table

Orbital Filling

4. Except for helium, the group _______ indicate the sum of electrons in the ns and np orbitals in the highest principal energy level that contains electrons (where n is the number that indicates a particular principal energy level). These electrons are the _______ electrons.
Electron Configurations and the Periodic Table

Orbital Filling

4. Except for helium, the group (NUMBERS) indicate the sum of electrons in the ns and np orbitals in the highest principal energy level that contains electrons (where n is the number that indicates a particular principal energy level). These electrons are the (VALENCE) electrons.
43. Electron Configurations and the Periodic Table

Exercise

Determine the expected electron configurations for each of the following.
a) S
b) Ba
c) Eu
Electron Configurations and the Periodic Table

Exercise

Determine the expected electron configurations for each of the following.
a) S
1s²2s²2p⁶3s²3p⁴ or [Ne]3s²3p⁴
b) Ba
[Xe]6s²
c) Eu
[Xe]6s²4f⁷
44.	Atomic Properties and the Periodic Table

Metals and Nonmetals

•	Metals tend to ____ electrons to form positive ions.
•	Nonmetals tend to ____ electrons to form negative ions.
44. Atomic Properties and the Periodic Table

Metals and Nonmetals

• Metals tend to ____ electrons to form positive ions.
• Nonmetals tend to ____ electrons to form negative ions.
Atomic Properties and the Periodic Table

Metals and Nonmetals

•	Metals tend to (LOSE) electrons to form positive ions.
•	Nonmetals tend to (GAIN) electrons to form negative ions.
Atomic Properties and the Periodic Table

Metals and Nonmetals

• Metals tend to (LOSE) electrons to form positive ions.
• Nonmetals tend to (GAIN) electrons to form negative ions.
45. Atomic Properties and the Periodic Table

Ionization Energy

• Energy required to ______ an electron from a gaseous atom or ion.
 X(g) → X⁺(g) + e⁻
Mg → Mg⁺ + e⁻ I₁ = 735 kJ/mol (1st IE)
Mg⁺ → Mg²⁺ + e⁻ I₂ = 1445 kJ/mol (2nd IE)
Mg²⁺ → Mg³⁺ + e⁻ I₃ = 7730 kJ/mol *(3rd IE)

*____ electrons are bound much more tightly than _______ electrons.
Atomic Properties and the Periodic Table

Ionization Energy

• Energy required to (REMOVE) an electron from a gaseous atom or ion.
 X(g) → X⁺(g) + e⁻
Mg → Mg⁺ + e⁻ I₁ = 735 kJ/mol (1st IE)
Mg⁺ → Mg²⁺ + e⁻ I₂ = 1445 kJ/mol (2nd IE)
Mg²⁺ → Mg³⁺ + e⁻ I₃ = 7730 kJ/mol *(3rd IE)

*(CORE) electrons are bound much more tightly than (VALENCE) electrons.
46. Atomic Properties and the Periodic Table

Ionization Energy

• In general, as we go across a period from left to right, the first __________ energy increases.
• Why?
 Electrons added in the same principal quantum level do not completely shield the increasing nuclear ______ caused by the added protons.
 Electrons in the same principal quantum level are generally more strongly _____ from left to right on the periodic table.
Atomic Properties and the Periodic Table

Ionization Energy

• In general, as we go across a period from left to right, the first (IONIZATION) energy increases.
• Why?
 Electrons added in the same principal quantum level do not completely shield the increasing nuclear (CHARGE) caused by the added protons.
 Electrons in the same principal quantum level are generally more strongly (BOUND) from left to right on the periodic table.
47.	Atomic Properties and the Periodic Table

__________ Energy
47. Atomic Properties and the Periodic Table

__________ Energy
Atomic Properties and the Periodic Table

(IONIZATION) Energy
Atomic Properties and the Periodic Table

(IONIZATION) Energy
48. Atomic Properties and the Periodic Table

Ionization Energy

• In general, as we go down a group from top to bottom, the first ionization energy _________.
• Why?
 The electrons being removed are, on average, farther from the _______.
Atomic Properties and the Periodic Table

Ionization Energy

• In general, as we go down a group from top to bottom, the first ionization energy (DECREASES).
• Why?
 The electrons being removed are, on average, farther from the (NUCLEUS).
49.	Atomic Properties and the Periodic Table

__________ Energy
49. Atomic Properties and the Periodic Table

__________ Energy
Atomic Properties and the Periodic Table

(IONIZATION) Energy
Atomic Properties and the Periodic Table

(IONIZATION) Energy
50. Atomic Properties and the Periodic Table

Concept Check

Which atom would require more energy to remove an electron? Why?

Na
Cl
Atomic Properties and the Periodic Table

Concept Check

Which atom would require more energy to remove an electron? Why?

Cl
51. Atomic Properties and the Periodic Table

Concept Check

Which atom would require more energy to remove an electron? Why?

Li
Cs
Atomic Properties and the Periodic Table

Concept Check

Which atom would require more energy to remove an electron? Why?

Li
52. Atomic Properties and the Periodic Table

Atomic Size

• In general as we go across a period from left to right, the atomic radius _________.
 Effective nuclear charge _________, therefore the valence electrons are drawn closer to the nucleus, decreasing the size of the atom.
• In general atomic radius increases in going down a group.
 _______ sizes increase in successive principal quantum levels.
Atomic Properties and the Periodic Table

Atomic Size

• In general as we go across a period from left to right, the atomic radius (DECREASES).
 Effective nuclear charge (INCREASES), therefore the valence electrons are drawn closer to the nucleus, decreasing the size of the atom.
• In general atomic radius increases in going down a group.
 (ORBITAL) sizes increase in successive principal quantum levels.
53.	Atomic Properties and the Periodic Table

Relative ______ _____ for Selected Atoms
53. Atomic Properties and the Periodic Table

Relative ______ _____ for Selected Atoms
Atomic Properties and the Periodic Table

Relative (ATOMIC SIZES) for Selected Atoms
Atomic Properties and the Periodic Table

Relative (ATOMIC SIZES) for Selected Atoms
54. Atomic Properties and the Periodic Table

Concept Check

Which should be the larger atom? Why?

Na
Cl
Atomic Properties and the Periodic Table

Concept Check

Which should be the larger atom? Why?

Na
55. Atomic Properties and the Periodic Table

Concept Check

Which should be the larger atom? Why?

Li
Cs
Atomic Properties and the Periodic Table

Concept Check

Which should be the larger atom? Why?

Cs
56. Atomic Properties and the Periodic Table

Concept Check

Which is larger?

• The hydrogen 1s orbital
• The lithium 1s orbital

Which is lower in energy?

• The hydrogen 1s orbital
• The lithium 1s orbital
Atomic Properties and the Periodic Table

Concept Check

Which is larger?

• The hydrogen 1s orbital

Which is lower in energy?

• The lithium 1s orbital
57. Atomic Properties and the Periodic Table

Exercise

Arrange the elements oxygen (O), fluorine (F), and sulfur (S) according to increasing:
 Ionization energy
 Atomic size
Atomic Properties and the Periodic Table

Exercise

Arrange the elements oxygen (O), fluorine (F), and sulfur (S) according to increasing:
 Ionization energy
S, O, F
 Atomic size
F, O, S
1. The concept of the nuclear atom left unanswered questions about

1. why the nucleus is so dense.
2. why the negative electrons do not collapse into the positive nucleus.
3. what isotopes are.
4. the mass of the nucleus.
Choice #2 properly explains that the nuclear atom did not explain the stable position of electrons relative to the positive nucleus.

Section 11.1: Rutherford’s Atom
2. The relationship between the wavelength and frequency of radiation is expressed by the equation:

1. λν = c
2. λ = 1/ν
3. ν = cλ
4. ν = 1/c + λ
Choice #1 provides the correct relationship:
Wavelength (λ) × frequency (ν) = the speed of light in a vacuum (c)

Section 11.2: Electromagnetic Radiation
3. The concept of photons emphasizes the particle-like nature of light rather than its wave-like nature. A photon of infrared light has more energy in it than a photon of

1. visible light
2. ultraviolet light
3. x-rays
4. microwave radiation
Choice #4. X-rays, UV light, and visible radiation are all more energetic than infrared.

Section 11.2: Electromagnetic Radiation
4. When the electrons in atoms receive energy from a source, they go into an excited state but quickly return back to its ground state by

1. nuclear fusion.
2. emitting a photon of light.
3. emitting a proton.
4. absorbing heat.
Choice #2 correctly describes a common mode of relaxation of excited state atoms.

Section 11.3: Emission of Energy by Atoms
5. The concept that electrons in an atom can only occupy specific energy levels is equivalent to saying that the energy levels are

1. varied.
2. quantized.
3. superimposed.
4. continuous.
Choice #2 reflects the concept that there are specific allowable energy levels in which electrons can exist.

Section 11.4: The Energy Levels of Hydrogen
6. Which of the following is true regarding the atom?

1. All atoms of the same element are identical.
2. As verified by Rutherford, only positively charged particles called protons are found inside the nucleus.
3. Electrons located further from the nucleus have more predictable behavior because they contain less energy.
4. Electrons display both particle-like behavior (they have mass) and wave-like behavior (they are associated with probability).
Choice #4 is the correct answer. Not all atoms of the same element are identical (due to isotopes and ions). Neutrons are also found in the nucleus of an atom in addition to protons. Electrons located further from the nucleus have less predictable behavior and contain more energy.

Section 11.4: The Energy Levels of Hydrogen
7. When an electron is excited in an atom or ion

1. only specific quantities of energy are released in order for the electron to return to its ground state.
2. white light is never observed when the electron returns to its ground state.
3. the electron is only excited to certain energy levels.
4. All of the above statements are true when an electron is excited.
Choice #4 is the correct answer.

Section 11.4: The Energy Levels of Hydrogen
8. How many of the following statements are true concerning an electron in its ground state?

I. The electron must be in its lowest-energy state.
II. Energy must be applied to the electron in order to excite it.
III. The electron must be located on the nucleus of an atom.
IV. The electron can release energy to obtain a lower ground state.
1. 1
2. 2
3. 3
4. 4
Choice #2 is the correct answer. I and II are correct.

Section 11.4: The Energy Levels of Hydrogen
9. The Bohr model of the hydrogen atom explained the emission spectrum of the hydrogen atom with the concept of

1. quantized energy levels.
2. quantized electron orbits.
3. transitions of electrons from higher to lower energy levels.
4. all of the above
Choice #4 reflects the quantized nature of the electron energy levels and the production of photons of specific energies that correspond to transitions between quantized levels.

Section 11.5: The Bohr Model of the Atom
10. The atomic model developed by Neils Bohr only works for the hydrogen atom. Choose the best answer that describes why this model does not work for other atoms.

1. The hydrogen atom has only one proton. The nuclei of other atoms would be able to pull the electrons out of their orbits.
2. The hydrogen atom has only one electron. Since other elements have more than one electron, the repulsions between electrons cannot be described by this model.
3. Other atoms are larger than hydrogen atoms, so they fill up too much space to be defined by orbits.
4. Hydrogen is the only element that has quantized energy levels like those described by this model.
Choice #2 is the correct answer.

Section 11.5: The Bohr Model of the Atom
11. In the wave mechanical model of the atom, electron orbitals are probability density diagrams of where the electron is likely to be found. These maps of where the electron is most likely to be were developed

1. from atomic emission spectra.
2. by Niels Bohr.
3. by experiments with a cathode ray tube.
4. from mathematical analyses by Schrödinger.
Choice #4 presents the idea that electron orbitals are mathematical descriptions of electron probability densities, derived from solutions to the Schrödinger equation.

Section 11.6: The Wave Mechanical Model of the Atom
12. At the third principal energy level of an atom, the number of orbitals equals

1. 3
2. 6
3. 9
4. 12
Choice #3 correctly identifies the sum of the orbitals at that level: one 3s orbital + three 3p orbitals + five 3d orbitals = nine orbitals.

Section 11.7: The Hydrogen Orbitals
13. The number of electrons that can be accommodated in the n = 2 level is:

1. 2
2. 4
3. 6
4. 8
Choice #4 correctly predicts that 8 electrons can fit in the n = 2 level. Two electrons fill the 2s orbital and 6 electrons fill the set of three 2p orbitals.

Section 11.8: The Wave Mechanical Model: Further Development
14. Which of the following statements best describes the movement of electrons in a p orbital?

1. 1.The electron movement cannot be exactly determined.
2. The electrons move within the two lobes of the p orbital, but never beyond the outside surface of the orbital.
3. The electrons are concentrated at the center (node) of the two lobes.
4. The electrons move along the outer surface of the p orbital, similar to a “figure 8” type of movement.
Choice #1 is the correct answer.

Section 11.8: The Wave Mechanical Model: Further Development
15. Write out the electron configuration for an atom of phosphorus and indicate the number of unpaired electrons

1. [Ne]3s²3p³ ; 3 unpaired electrons
2. [Ne]2s²2p⁶3s²3p³ ; 3 unpaired electrons
3. [Ne]3s²3p³ ; 5 unpaired electrons
4. [Ne]3p⁵ ; 5 unpaired electrons
Choice #1 is correct. Each of the three 3p orbitals has one unpaired electron in it. The electrons in the 3s orbital are paired.

Section 11.9: Electron Arrangements in the First Eighteen Atoms on the Periodic Table
16. The number of valence electrons in an atom of Cl is:

1. 7
2. 5
3. 3
4. 1
Choice #1 is correct. Cl has 2 valence electrons in the 3s orbital and 5 in the 3p subset.

Section 11.9: Electron Arrangements in the First Eighteen Elements on the Periodic Table
17. How many unpaired electrons does sulfur contain in its ground state?

1. 0
2. 1
3. 2
4. 3
Choice #3 is correct. There are 4 electrons in the p orbitals, with one of the p orbitals containing paired electrons and two of the p orbitals containing unpaired electrons.

Section 11.9: Electron Arrangements in the First Eighteen Elements on the Periodic Table
18. How many unpaired electrons does the element cobalt (Co) have in its lowest energy state?

1. 0
2. 2
3. 3
4. 7
Choice #3 is correct. There are 7 electrons in the d orbitals, with two of the d orbitals containing paired electrons and three of the d orbitals containing unpaired electrons.

Section 11.9: Electron Arrangements in the First Eighteen Elements on the Periodic Table
19. Can an electron in a phosphorus atom ever be in a 3d orbital? Choose the best answer.

1.Yes. An electron can be excited into a 3d orbital.
2. Yes. A ground-state electron in phosphorus is located in a 3d orbital.
3. No. Only transition metal atoms can have electrons located in the d orbitals.
4. No. This would not correspond to phosphorus’ electron arrangement in its ground state.
Choice #1 is the correct answer.

Section 11.9: Electron Arrangements in the First Eighteen Elements on the Periodic Table
20. The Lanthanides correspond to filling up the ____set of orbitals.

1. 3d
2. 3f
3. 4f
4. 5f
Choice #3 correctly associates the 14 elements that make up the lanthanides with 14 available slots for electrons in the set of seven 4f orbitals.

Section 11.10: Electron Configurations and the Periodic Table
21. Determine the expected electron configuration for the element Eu.

1. [Xe]6s²4f⁶
2. [Rn]7s²5f⁷
3. [Xe]4f⁷
4. [Xe]6s²4f⁷
Choice #4 is the correct answer.

Section 11.10: Electron Configurations and the Periodic Table
22. Which element would be likely to have the lowest ionization energy?

1. Na
2. Al
3. Cl
4. Cs
Choice #4 should be selected. Ionization energy increases from left to right, but decreases from top to bottom.

Section 11.11: Atomic Properties and the Periodic Table
23. Which of the following elements has the largest ionization energy?

1. P
2. Al
3. Cl
4. K
Choice #3 is the correct answer. Ionization energy increases from left to right, but decreases from top to bottom.

Section 11.11: Atomic Properties and the Periodic Table
24. Which of the following elements has the largest ionization energy?

1. S
2. Ba
3. Cr
4. Mg
Choice #1 is the correct answer. Ionization energy increases from left to right, but decreases from top to bottom.

Section 11.11: Atomic Properties and the Periodic Table
25. Which of the following atoms is likely to have the largest atomic radius?

1. Na
2. Al
3. Cl
4. K
Choice #4. Atomic radius decreases as one proceeds across the third period, but increases from top to bottom.

Section 11.11: Atomic Properties and the Periodic Table
26. Which of the following atoms is likely to have the largest atomic radius?

1. Ca
2. Sr
3. N
4. Al
Choice #2. Atomic radius decreases as one proceeds across the third period, but increases from top to bottom.

Section 11.11: Atomic Properties and the Periodic Table
27. Which of the following is ranked in order of largest to smallest atomic radius?

1. F > S > Ge > Mn > Rb
2. Mn > Rb > F > S > Ge
3. Rb > Mn > Ge > S > F
4. Rb > Ge > Mn > F > S
Choice #3. Atomic radius decreases as one proceeds across the third period, but increases from top to bottom.

Section 11.11: Atomic Properties and the Periodic Table
28. Rank the following from smallest to largest atomic radius.

1. O, Zn, Ca, Ba
2. O, Ca, Zn, Ba
3. Ba, Ca, Zn, O
4. O, Zn, Ba, Ca
Choice #1. Atomic radius decreases as one proceeds across the third period, but increases from top to bottom.

Section 11.11: Atomic Properties and the Periodic Table
29. Which has a smaller atomic radius, an atom of fluorine or an atom of lithium? Choose the best answer.

1. They are the same size because their electrons are contained in the same principle energy level.
2. An atom of lithium is smaller than an atom of fluorine because lithium has fewer protons.
3. An atom of fluorine is smaller than an atom of lithium because with more protons the electrons are pulled closer to the nucleus.
4. An atom of fluorine is smaller than an atom of lithium because fluorine has a high ionization energy and will not lose electrons as easily as lithium.
Choice #3 is the correct answer.

Section 11.11: Atomic Properties and the Periodic Table
30. Which of the following statements is true?

1. The krypton 1s orbital is larger than the helium 1s orbital because krypton contains more electrons.
2. The krypton 1s orbital is smaller than the helium 1s orbital because krypton’s nuclear charge draws the electrons closer.
3. The krypton 1s orbital and helium 1s orbital are the same size because both s orbitals can only have two electrons.
4. The krypton 1s orbital is larger than the helium 1s orbital because krypton’s ionization energy is lower so it’s easier to remove electrons.
Choice #2 is the correct answer.

Section 11.11: Atomic Properties and the Periodic Table
The force that holds two atoms together in a compound.
1. Bond
The energy required to break a given chemical bond.
2. Bond Energy
The attraction between oppositely charged ions.
3. Ionic Bonding
A compound that results when a metal reacts with a nonmetal to form cations and anions.
4. Ionic Compound
A type of bonding in which atoms share electrons.
5. Covalent Bonding
A covalent bond in which the electrons are not shared equally because one atom attracts them more strongly than the other.
6. Polar Covalent Bond
The tendency of an atom in a molecule to attract shared electrons to itself.
7. Electronegativity
A property of a molecule whereby the charge distribution can be represented by a center of positive charge and a center of negative charge.
8. Dipole Moment
A diagram of a molecule showing how the valence electrons are arranged among the atoms in the molecule.
9. Lewis Structure
The distribution of valence electrons when hydrogen atoms—which end up with only two valence electrons—experience chemical bonding with other atoms; most other elements follow the octet rule (online definition).
10. Duet Rule
The observation that atoms of nonmetals form the most stable molecules when they are surrounded by eight electrons (to fill their valence orbitals).
11. Octet Rule
An electron pair found in the space between two atoms.
12. Bonding Pair
An electron pair that is localized on a given atom; an electron pair not involved in bonding.
13. Lone (Unshared) Pair
A bond in which two atoms share one pair of electrons.
14. Single Bond
A bond in which two atoms share two pairs of electrons.
15. Double Bond
A bond in which two atoms share three pairs of electrons.
16. Triple Bond
A condition occurring when more than one valid Lewis structure can be written for a particular molecule; the actual electron structure is represented not by any one of the Lewis structures but by the average of all of them.
17. Resonance
Various Lewis structures.
18. Resonance Structure
The three-dimensional arrangement of atoms in a molecule.
19. Molecular (Geometric) Structure
The angle between bonds sharing a common atom; also known as valence angle (online definition).
20. Bond Angle
A plot structure that runs in a chronological or logical cause-and-effect sequence (online definition).
21. Linear Structure
The geometric arrangement formed when three electron groups maximize their separation around a central atom (online definition).
22. Trigonal Planar Structure
a
23. Tetrahedral Structure
A model the main postulate of which is that the structure around a given atom in a molecule is determined principally by the tendency to minimize electron-pair repulsions.
24. Valence Shell Electron Pair Repulsion (VSEPR) Model
The geometric arrangement formed when four electron groups maximize their separation around a central atom; when all four groups are bonding groups, the molecular shape is ___________ (online definition).
25. Tetrahedral Arrangement
A molecular geometry with one atom at the apex and three atoms at the corners of a trigonal base (online definition).
26. Trigonal Pyramid
1. Questions to Consider

• What is meant by the term “chemical bond”?
• Why do atoms bond with each other to form compounds?
• How do atoms bond with each other to form compounds?
Questions to Consider

• What is meant by the term “chemical bond”?
• Why do atoms bond with each other to form compounds?
• How do atoms bond with each other to form compounds?
2.	Types of Chemical Bonds

A Chemical Bond

•	No simple, and yet complete, way to define this.
•	Forces that hold groups of _____ together and make them function as a unit.
•	A bond will form if the energy of the aggregate is _____ than that of the
2. Types of Chemical Bonds

A Chemical Bond

• No simple, and yet complete, way to define this.
• Forces that hold groups of _____ together and make them function as a unit.
• A bond will form if the energy of the aggregate is _____ than that of the separated atoms.
• Bond energy – energy required to break a ________ ____.
Types of Chemical Bonds

A Chemical Bond

•	No simple, and yet complete, way to define this.
•	Forces that hold groups of (ATOMS) together and make them function as a unit.
•	A bond will form if the energy of the aggregate is (LOWER) than that of th
Types of Chemical Bonds

A Chemical Bond

• No simple, and yet complete, way to define this.
• Forces that hold groups of (ATOMS) together and make them function as a unit.
• A bond will form if the energy of the aggregate is (LOWER) than that of the separated atoms.
• Bond energy – energy required to break a (CHEMICAL BOND).
3.	Types of Chemical Bonds

Ionic Bonding

•	Ionic compound results when a _____ reacts with a ________. 
•	_________ are transferred.
3. Types of Chemical Bonds

Ionic Bonding

• Ionic compound results when a _____ reacts with a ________.
• _________ are transferred.
Types of Chemical Bonds

Ionic Bonding

•	Ionic compound results when a (METAL) reacts with a (NONMETAL). 
•	(ELECTRONS) are transferred.
Types of Chemical Bonds

Ionic Bonding

• Ionic compound results when a (METAL) reacts with a (NONMETAL).
• (ELECTRONS) are transferred.
4.	Types of Chemical Bonds

Covalent Bonding

•	A covalent bond results when _________ are shared by ______.
4. Types of Chemical Bonds

Covalent Bonding

• A covalent bond results when _________ are shared by ______.
Types of Chemical Bonds

Covalent Bonding

•	A covalent bond results when (ELECTRONS) are shared by (NUCLEI).
Types of Chemical Bonds

Covalent Bonding

• A covalent bond results when (ELECTRONS) are shared by (NUCLEI).
5.	Types of Chemical Bonds

Polar Covalent Bond

•	_______ sharing of electrons between atoms in a molecule.
•	One atom attracts the _________ more than the other atom.
•	Results in a ______ separation in the bond (partial positive and partial negat
5. Types of Chemical Bonds

Polar Covalent Bond

• _______ sharing of electrons between atoms in a molecule.
• One atom attracts the _________ more than the other atom.
• Results in a ______ separation in the bond (partial positive and partial negative charge).
Types of Chemical Bonds

Polar Covalent Bond

•	(UNEQUAL) sharing of electrons between atoms in a molecule.
•	One atom attracts the (ELECTRONS) more than the other atom.
•	Results in a (CHARGE) separation in the bond (partial positive and partial ne
Types of Chemical Bonds

Polar Covalent Bond

• (UNEQUAL) sharing of electrons between atoms in a molecule.
• One atom attracts the (ELECTRONS) more than the other atom.
• Results in a (CHARGE) separation in the bond (partial positive and partial negative charge).
6. Types of Chemical Bonds

Concept Check

What is meant by the term “chemical bond?”

Why do atoms bond with each other to form molecules?

How do atoms bond with each other to form molecules?
Types of Chemical Bonds

Concept Check

What is meant by the term “chemical bond?”

Why do atoms bond with each other to form molecules?

How do atoms bond with each other to form molecules?
7. Electronegativity

• The ability of an atom in a molecule to _______ shared electrons to itself.
• For a molecule HX, the relative electronegativities of the H and X atoms are determined by comparing the ________ H–X bond energy with the “expected” H–X bond energy.
Electronegativity

• The ability of an atom in a molecule to (ATTRACT) shared electrons to itself.
• For a molecule HX, the relative electronegativities of the H and X atoms are determined by comparing the (MEASURED) H–X bond energy with the “expected” H–X bond energy.
8. Electronegativity

• On the periodic table, electronegativity generally _________ across a period and _________ down a group.
• The range of electronegativity values is from 4.0 for fluorine (the ____ electronegative) to 0.7 for cesium and francium (the _____ electronegative).
Electronegativity

• On the periodic table, electronegativity generally (INCREASES) across a period and (DECREASES) down a group.
• The range of electronegativity values is from 4.0 for fluorine (the [MOST] electronegative) to 0.7 for cesium and francium (the [LEAST] electronegative).
9.	Electronegativity

Electronegativity Values for Selected ________
9. Electronegativity

Electronegativity Values for Selected ________
Electronegativity

Electronegativity Values for Selected (ELEMENTS)
Electronegativity

Electronegativity Values for Selected (ELEMENTS)
10. Electronegativity

Concept Check

If lithium and fluorine react, which has more attraction for an electron? Why?

In a bond between fluorine and iodine, which has more attraction for an electron? Why?
Electronegativity

Concept Check

If lithium and fluorine react, which has more attraction for an electron? Why?

In a bond between fluorine and iodine, which has more attraction for an electron? Why?
11. Electronegativity

Concept Check

What is the general trend for electronegativity across rows and down columns on the periodic table?

Explain the trend.
Electronegativity

Concept Check

What is the general trend for electronegativity across rows and down columns on the periodic table?

Explain the trend.
12.	Electronegativity

•	The ________ of a bond depends on the difference between the electronegativity values of the atoms forming the bond.
12. Electronegativity

• The ________ of a bond depends on the difference between the electronegativity values of the atoms forming the bond.
Electronegativity

•	The (POLARITY) of a bond depends on the difference between the electronegativity values of the atoms forming the bond.
Electronegativity

• The (POLARITY) of a bond depends on the difference between the electronegativity values of the atoms forming the bond.
13.	Electronegativity

The Relationship Between Electronegativity and ____ ____
13. Electronegativity

The Relationship Between Electronegativity and ____ ____
Electronegativity

The Relationship Between Electronegativity and (BOND TYPE)
Electronegativity

The Relationship Between Electronegativity and (BOND TYPE)
14. Electronegativity

Exercise

Arrange the following bonds from most to least polar:

a) N–F, O–F, C–F
b) C–F, N–O, Si–F
c) Cl–Cl, B–Cl, S–Cl
Electronegativity

Exercise

Arrange the following bonds from most to least polar:

a) N–F, O–F, C–F
b) C–F, N–O, Si–F
c) Cl–Cl, B–Cl, S–Cl

a) C–F, N–F, O–F
b) Si–F, C–F, N–O
c) B–Cl, S–Cl, Cl–Cl
15. Electronegativity

Concept Check

Which of the following bonds would be the least polar yet still be considered polar covalent?

Mg–O, C–O, O–O, Si–O, N–O
Electronegativity

Concept Check

Which of the following bonds would be the least polar yet still be considered polar covalent?

Mg–O, C–O, O–O, Si–O, N–O

N-O
16. Electronegativity

Concept Check

Which of the following bonds would be the most polar without being considered ionic?

Mg–O, C–O, O–O, Si–O, N–O
Electronegativity

Concept Check

Which of the following bonds would be the most polar without being considered ionic?

Mg–O, C–O, O–O, Si–O, N–O
17.	Bond Polarity and Dipole Moments

Dipole Moment

•	Property of a ________ whose charge distribution can be represented by a center of positive charge and a center of negative charge.
•	Use an _____ to represent a dipole moment.
	Point to the ne
17. Bond Polarity and Dipole Moments

Dipole Moment

• Property of a ________ whose charge distribution can be represented by a center of positive charge and a center of negative charge.
• Use an _____ to represent a dipole moment.
 Point to the negative charge center with the tail of the arrow indicating the positive center of charge.
Bond Polarity and Dipole Moments

Dipole Moment

•	Property of a (MOLECULE) whose charge distribution can be represented by a center of positive charge and a center of negative charge.
•	Use an (ARROW) to represent a dipole moment.
	Point to the ne
Bond Polarity and Dipole Moments

Dipole Moment

• Property of a (MOLECULE) whose charge distribution can be represented by a center of positive charge and a center of negative charge.
• Use an (ARROW) to represent a dipole moment.
 Point to the negative charge center with the tail of the arrow indicating the positive center of charge.
18.	Bond Polarity and Dipole Moments

Dipole Moment in a _____ Molecule
18. Bond Polarity and Dipole Moments

Dipole Moment in a _____ Molecule
Bond Polarity and Dipole Moments

Dipole Moment in a (WATER) Molecule
Bond Polarity and Dipole Moments

Dipole Moment in a (WATER) Molecule
19.	Bond Polarity and Dipole Moments

•	The ________ of water affects its properties. 
	Permits _____ compounds to dissolve in it. 
	Causes water to remain liquid at higher temperature.
19. Bond Polarity and Dipole Moments

• The ________ of water affects its properties.
 Permits _____ compounds to dissolve in it.
 Causes water to remain liquid at higher temperature.
Bond Polarity and Dipole Moments

•	The (POLARITY) of water affects its properties. 
	Permits (IONIC) compounds to dissolve in it. 
	Causes water to remain liquid at higher temperature.
Bond Polarity and Dipole Moments

• The (POLARITY) of water affects its properties.
 Permits (IONIC) compounds to dissolve in it.
 Causes water to remain liquid at higher temperature.
20. Stable Electron Configurations and Charges on Ions

Review

• Group 1 metals always form __ cations.
• Group 2 metals always form __ cations.
• Aluminum in Group 3 always forms a __ cation.
• Group 6 elements always form __ anions.
• Group 7 nonmetals form __ anions.
Stable Electron Configurations and Charges on Ions

Review

• Group 1 metals always form (1+) cations.
• Group 2 metals always form (2+) cations.
• Aluminum in Group 3 always forms a (3+) cation.
• Group 6 elements always form (2–) anions.
• Group 7 nonmetals form (1–) anions.
21. Stable Electron Configurations and Charges on Ions

Stable Compounds

• Atoms in ______ compounds usually have a noble gas electron configuration.
 ______ lose electrons to reach a noble gas configuration.
 _________ gain electrons to reach a noble gas configuration.
Stable Electron Configurations and Charges on Ions

Stable Compounds

• Atoms in (STABLE) compounds usually have a noble gas electron configuration.
 (METALS) lose electrons to reach a noble gas configuration.
 (NONMETALS) gain electrons to reach a noble gas configuration.
22.	Stable Electron Configurations and Charges on Ions

The Formation of ____ by Metals and Nonmetals
22. Stable Electron Configurations and Charges on Ions

The Formation of ____ by Metals and Nonmetals
Stable Electron Configurations and Charges on Ions

The Formation of (IONS) by Metals and Nonmetals
Stable Electron Configurations and Charges on Ions

The Formation of (IONS) by Metals and Nonmetals
23. Stable Configurations and Charges on Ions

Electron Configurations of Ions

1. Representative (main-group) metals form ions by ______ enough electrons to achieve the configuration of the previous noble gas.
2. Nonmetals form ions by _______ enough electrons to achieve the configuration of the next noble gas.
Stable Configurations and Charges on Ions

Electron Configurations of Ions

1. Representative (main-group) metals form ions by (LOSING) enough electrons to achieve the configuration of the previous noble gas.
2. Nonmetals form ions by (GAINING) enough electrons to achieve the configuration of the next noble gas.
24. Stable Electron Configurations and Charges on Ions

Electron Configurations and Bonding

1. When a nonmetal and a Group 1, 2, or 3 metal react to form a binary ionic compound, the ions form so that the valence-electron configuration of the nonmetal ________ the electron configuration of the next noble gas atom. The valence orbitals of the metal are _______ to achieve the configuration of the previous noble gas.
2. When two nonmetals react to form a covalent bond, they _____ electrons in a way that completes the valence-electron configurations of both atoms.
Stable Electron Configurations and Charges on Ions

Electron Configurations and Bonding

1. When a nonmetal and a Group 1, 2, or 3 metal react to form a binary ionic compound, the ions form so that the valence-electron configuration of the nonmetal (ACHIEVES) the electron configuration of the next noble gas atom. The valence orbitals of the metal are (EMPTIED) to achieve the configuration of the previous noble gas.
2. When two nonmetals react to form a covalent bond, they (SHARE) electrons in a way that completes the valence-electron configurations of both atoms.
25.	Stable Electron Configurations and Charges on Ions

Predicting Formulas of Ionic Compounds

•	Chemical compounds are always electrically _______.
25. Stable Electron Configurations and Charges on Ions

Predicting Formulas of Ionic Compounds

• Chemical compounds are always electrically _______.
Stable Electron Configurations and Charges on Ions

Predicting Formulas of Ionic Compounds

•	Chemical compounds are always electrically (NEUTRAL).
Stable Electron Configurations and Charges on Ions

Predicting Formulas of Ionic Compounds

• Chemical compounds are always electrically (NEUTRAL).
26. Stable Electron Configurations and Charges on Ions

Concept Check

What is the expected ground state electron configuration for Te2–?

a) [Kr]5s²4d¹⁰5p⁴
b) [Kr]5s²4d¹⁰4f¹⁴5p⁶
c) [Kr]5s²4d¹⁰5p⁶
d) [Ar]5s²4d¹⁰5p²
Stable Electron Configurations and Charges on Ions

Concept Check

What is the expected ground state electron configuration for Te2–?

c) [Kr]5s²4d¹⁰5p⁶
27. Stable Electron Configurations and Charges on Ions

Concept Check

What is the correct electron configuration for the most stable form of the sulfur ion in an ionic compound?

a) 1s²2s²2p⁶3s²
b) 1s²2s²2p⁶3s²3p²
c) 1s²2s²2p⁶3s²3p⁴
d) 1s²2s²2p⁶3s²3p⁶
Stable Electron Configurations and Charges on Ions

Concept Check

What is the correct electron configuration for the most stable form of the sulfur ion in an ionic compound?

d) 1s²2s²2p⁶3s²3p⁶
28.	Ionic Bonding and Structures of Ionic Compounds

Structures of Ionic Compounds

•	Ions are packed together to ________ the attractions between ions.
28. Ionic Bonding and Structures of Ionic Compounds

Structures of Ionic Compounds

• Ions are packed together to ________ the attractions between ions.
Ionic Bonding and Structures of Ionic Compounds

Structures of Ionic Compounds

•	Ions are packed together to (MAXIMIZE) the attractions between ions.
Ionic Bonding and Structures of Ionic Compounds

Structures of Ionic Compounds

• Ions are packed together to (MAXIMIZE) the attractions between ions.
29.	Ionic Bonding and Structures of Ionic Compounds

Structures of Ionic Compounds

•	_______ are always smaller than the parent atom.
•	______ are always larger than the parent atom.
29. Ionic Bonding and Structures of Ionic Compounds

Structures of Ionic Compounds

• _______ are always smaller than the parent atom.
• ______ are always larger than the parent atom.
Ionic Bonding and Structures of Ionic Compounds

Structures of Ionic Compounds

•	(CATIONS) are always smaller than the parent atom.
•	(ANIONS) are always larger than the parent atom.
Ionic Bonding and Structures of Ionic Compounds

Structures of Ionic Compounds

• (CATIONS) are always smaller than the parent atom.
• (ANIONS) are always larger than the parent atom.
30. Ionic Bonding and Structures of Ionic Compounds

Isoelectronic Series

• A series of ____/_____ containing the same number of electrons.
O²⁻, F⁻, Ne, Na⁺, Mg²⁺, and Al³⁺
Ionic Bonding and Structures of Ionic Compounds

Isoelectronic Series

• A series of (IONS/ATOMS) containing the same number of electrons.
O²⁻, F⁻, Ne, Na⁺, Mg²⁺, and Al³⁺
31. Ionic Bonding and Structures of Ionic Compounds

Concept Check

Choose an alkali metal, an alkaline earth metal, a noble gas, and a halogen so that they constitute an isoelectronic series when the metals and halogen are written as their most stable ions.
 What is the electron configuration for each species?
 Determine the number of electrons for each species.
 Determine the number of protons for each species.
 Rank the species according to increasing radius.
 Rank the species according to increasing ionization energy.
Ionic Bonding and Structures of Ionic Compounds

Concept Check

Choose an alkali metal, an alkaline earth metal, a noble gas, and a halogen so that they constitute an isoelectronic series when the metals and halogen are written as their most stable ions.
 What is the electron configuration for each species?
 Determine the number of electrons for each species.
 Determine the number of protons for each species.
 Rank the species according to increasing radius.
 Rank the species according to increasing ionization energy.
32. Ionic Bonding and Structures of Ionic Compounds

Concept Check

Rank the following from smallest to largest atomic radius:

Ar, S²⁻, Ca²⁺, K⁺, Cl⁻

a) Ar < K⁺ < Ca²⁺ < S²⁻ < Cl⁻
b) Ca²⁺ < K⁺ < Ar < Cl⁻ < S²⁻
c) Ar < Cl⁻ < S²⁻ < Ca²⁺ < K⁺
d) S²⁻ < Cl⁻ < Ar < K⁺ < Ca²⁺
Ionic Bonding and Structures of Ionic Compounds

Concept Check

Rank the following from smallest to largest atomic radius:

Ar, S²⁻, Ca²⁺, K⁺, Cl⁻

b) Ca²⁺ < K⁺ < Ar < Cl⁻ < S²⁻
33. Ionic Bonding and Structures of Ionic Compounds

Concept Check

Which atom or ion has the smallest radius?

a) O²⁺
b) O⁺
c) O
d) O²⁻
Ionic Bonding and Structures of Ionic Compounds

Concept Check

Which atom or ion has the smallest radius?

a) O²⁺
34. Ionic Bonding and Structures of Ionic Compounds

Ionic Compounds Containing Polyatomic Ions

• Polyatomic ions work in the same way as simple ions.
 The ________ _____ hold the polyatomic ion together so it behaves as a unit.
Ionic Bonding and Structures of Ionic Compounds

Ionic Compounds Containing Polyatomic Ions

• Polyatomic ions work in the same way as simple ions.
 The (COVALENT BONDS) hold the polyatomic ion together so it behaves as a unit.
35.	Lewis Structures

Lewis Structure

•	Shows how _______ electrons are arranged among atoms in a molecule.
•	Most important requirement
	Atoms achieve _____ ___ electron configuration (octet rule, duet rule).
35. Lewis Structures

Lewis Structure

• Shows how _______ electrons are arranged among atoms in a molecule.
• Most important requirement
 Atoms achieve _____ ___ electron configuration (octet rule, duet rule).
Lewis Structures

Lewis Structure

•	Shows how (VALENCE) electrons are arranged among atoms in a molecule.
•	Most important requirement
	Atoms achieve (NOBLE GAS) electron configuration (octet rule, duet rule).
Lewis Structures

Lewis Structure

• Shows how (VALENCE) electrons are arranged among atoms in a molecule.
• Most important requirement
 Atoms achieve (NOBLE GAS) electron configuration (octet rule, duet rule).
36.	Lewis Structures

Writing Lewis Structures

•	_______ _____ are shared between 2 atoms.
•	________ _____ (lone pairs) are not shared and not involved in bonding.
36. Lewis Structures

Writing Lewis Structures

• _______ _____ are shared between 2 atoms.
• ________ _____ (lone pairs) are not shared and not involved in bonding.
Lewis Structures

Writing Lewis Structures

•	(BONDING PAIRS) are shared between 2 atoms.
•	(UNSHARED PAIRS) (lone pairs) are not shared and not involved in bonding.
Lewis Structures

Writing Lewis Structures

• (BONDING PAIRS) are shared between 2 atoms.
• (UNSHARED PAIRS) (lone pairs) are not shared and not involved in bonding.
37. Lewis Structures

Steps for Writing Lewis Structures

1. Sum the _______ electrons from all the atoms.
2. Use a pair of electrons to form a ____ between each pair of bound atoms.
3. Atoms usually have _____ ___ configurations. Arrange the remaining electrons to satisfy the octet rule (or duet rule for hydrogen).
Lewis Structures

Steps for Writing Lewis Structures

1. Sum the (VALENCE) electrons from all the atoms.
2. Use a pair of electrons to form a (BOND) between each pair of bound atoms.
3. Atoms usually have (NOBLE GAS) configurations. Arrange the remaining electrons to satisfy the octet rule (or duet rule for hydrogen).
38. Lewis Structures

Steps for Writing Lewis Structures

1. Sum the valence electrons from ___ the atoms. (Use the periodic table.)

Example:

H₂O

2 (1 e⁻) + 6 e⁻ = 8 e⁻ total
Lewis Structures

Steps for Writing Lewis Structures

1. Sum the valence electrons from (ALL) the atoms. (Use the periodic table.)

Example:

H₂O

2 (1 e⁻) + 6 e⁻ = 8 e⁻ total
39. Lewis Structures

Steps for Writing Lewis Structures

2. Use a pair of electrons to form a ____ between each pair of bound atoms.

Example:

H₂O

H-O-H
Lewis Structures

Steps for Writing Lewis Structures

2. Use a pair of electrons to form a (BOND) between each pair of bound atoms.

Example:

H₂O

H-O-H
40.	Lewis Structures

Steps for Writing Lewis Structures

3.	Atoms usually have _____ ___ configurations. Arrange the remaining electrons to satisfy the octet rule (or duet rule for hydrogen).

Examples:

H2O and PBr3
40. Lewis Structures

Steps for Writing Lewis Structures

3. Atoms usually have _____ ___ configurations. Arrange the remaining electrons to satisfy the octet rule (or duet rule for hydrogen).

Examples:

H₂O and PBr₃
Lewis Structures

Steps for Writing Lewis Structures

3.	Atoms usually have (NOBLE GAS) configurations. Arrange the remaining electrons to satisfy the octet rule (or duet rule for hydrogen).

Examples:

H2O and PBr3
Lewis Structures

Steps for Writing Lewis Structures

3. Atoms usually have (NOBLE GAS) configurations. Arrange the remaining electrons to satisfy the octet rule (or duet rule for hydrogen).

Examples:

H₂O and PBr₃
41. Lewis Structures

Concept Check

Draw a Lewis structure for each of the following molecules:

a) H₂
b) F₂
c) HF
d) CH₄
Lewis Structures

Concept Check

Draw a Lewis structure for each of the following molecules:

a)	H2
b)	F2
c)	HF
d)	CH4
Lewis Structures

Concept Check

Draw a Lewis structure for each of the following molecules:

a) H₂
b) F₂
c) HF
d) CH₄
42. Lewis Structures of Molecules with Multiple Bonds

• Single bond – covalent bond in which 1 pair of electrons is shared by 2 atoms.
• Double bond – covalent bond in which 2 pairs of electrons are shared by 2 atoms.
• Triple bond – covalent bond in which 3 pairs of electrons are shared by 2 atoms.
Lewis Structures of Molecules with Multiple Bonds

•	Single bond – covalent bond in which 1 pair of electrons is shared by 2 atoms.
H–H
•	Double bond – covalent bond in which 2 pairs of electrons are shared by 2 atoms.
O=C=O
•	Triple bond – covalent
Lewis Structures of Molecules with Multiple Bonds

• Single bond – covalent bond in which 1 pair of electrons is shared by 2 atoms.
H–H
• Double bond – covalent bond in which 2 pairs of electrons are shared by 2 atoms.
O=C=O
• Triple bond – covalent bond in which 3 pairs of electrons are shared by 2 atoms.
43.	Lewis Structures of Molecules with Multiple Bonds

Resonance

•	A molecule shows resonance when more than one Lewis structure can be drawn for the ________.

NO3– = 24e–
43. Lewis Structures of Molecules with Multiple Bonds

Resonance

• A molecule shows resonance when more than one Lewis structure can be drawn for the ________.

NO₃⁻ = 24e⁻
Lewis Structures of Molecules with Multiple Bonds

Resonance

•	A molecule shows resonance when more than one Lewis structure can be drawn for the (MOLECULE).

NO3– = 24e–
Lewis Structures of Molecules with Multiple Bonds

Resonance

• A molecule shows resonance when more than one Lewis structure can be drawn for the (MOLECULE).

NO₃⁻ = 24e⁻
44.	Lewis Structures of Molecules with Multiple Bonds

Some Exceptions to the Octet Rule

•	Boron tends to form compounds in which the boron atom has _____ than eight electrons around it (it does not have a complete octet).

BH3 = 6e–

•	Molecules
44. Lewis Structures of Molecules with Multiple Bonds

Some Exceptions to the Octet Rule

• Boron tends to form compounds in which the boron atom has _____ than eight electrons around it (it does not have a complete octet).

BH₃ = 6e⁻

• Molecules containing ___ numbers of electrons like NO and NO₂.
Lewis Structures of Molecules with Multiple Bonds

Some Exceptions to the Octet Rule

•	Boron tends to form compounds in which the boron atom has (FEWER) than eight electrons around it (it does not have a complete octet).

BH3 = 6e–

•	Molecules c
Lewis Structures of Molecules with Multiple Bonds

Some Exceptions to the Octet Rule

• Boron tends to form compounds in which the boron atom has (FEWER) than eight electrons around it (it does not have a complete octet).

BH₃ = 6e⁻

• Molecules containing (ODD) numbers of electrons like NO and NO₂.
45. Lewis Structures of Molecules with Multiple Bonds

Concept Check

Draw a Lewis structure for each of the following molecules:

a) BF₃
b) CO₂
c) CCl₄
d) CN⁻
Lewis Structures of Molecules with Multiple Bonds

Concept Check

Draw a Lewis structure for each of the following molecules:

e)	BF3
f)	CO2
g)	CCl4
h)	CN–
Lewis Structures of Molecules with Multiple Bonds

Concept Check

Draw a Lewis structure for each of the following molecules:

a) BF₃
b) CO₂
c) CCl₄
d) CN⁻
46. Lewis Structures of Molecules with Multiple Bonds

Concept Check

Consider the following compounds:

CO₂, N₂, CCl₄

Which compound exhibits resonance?

a) CO₂
b) N₂
c) CCl₄
d) At least two of the above compounds exhibit resonance.
Lewis Structures of Molecules with Multiple Bonds

Concept Check

Consider the following compounds:

CO₂, N₂, CCl₄

Which compound exhibits resonance?

a) CO₂
47. Lewis Structures of Molecules with Multiple Bonds

Concept Check

Which of the following supports why Lewis structures are not a completely accurate way to draw molecules?

a) We cannot say for certain where an electron is located yet when drawing Lewis structures, we assume the electrons are right where we place them.
b) When adding up the number of valence electrons for a molecule, it is possible to get an odd number which would make it impossible to satisfy the octet rule for all atoms.
c) Both statements 1 and 2 above support why Lewis structures are not a completely accurate way to draw molecules.
d) Lewis structures are the most accurate way to draw molecules and are completely correct.
Lewis Structures of Molecules with Multiple Bonds

Concept Check

Which of the following supports why Lewis structures are not a completely accurate way to draw molecules?

c) Both statements 1 and 2 above support why Lewis structures are not a completely accurate way to draw molecules.
48.	Molecular Structure

•	Three dimensional arrangement of the _____ in a molecule.
48. Molecular Structure

• Three dimensional arrangement of the _____ in a molecule.
Molecular Structure

•	Three dimensional arrangement of the (ATOMS) in a molecule.
Molecular Structure

• Three dimensional arrangement of the (ATOMS) in a molecule.
49.	Molecular Structure

•	Linear structure – atoms in a ____
	Carbon dioxide
49. Molecular Structure

• Linear structure – atoms in a ____
 Carbon dioxide
Molecular Structure

•	Linear structure – atoms in a (LINE)
	Carbon dioxide
Molecular Structure

• Linear structure – atoms in a (LINE)
 Carbon dioxide
50.	Molecular Structure

•	Trigonal planar – atoms in a ________
	Boron trifluoride
50. Molecular Structure

• Trigonal planar – atoms in a ________
 Boron trifluoride
Molecular Structure

•	Trigonal planar – atoms in a (TRIANGLE)
	Boron trifluoride
Molecular Structure

• Trigonal planar – atoms in a (TRIANGLE)
 Boron trifluoride
51.	Molecular Structure

•	___________ Structure
	Methane
51. Molecular Structure

• ___________ Structure
 Methane
Molecular Structure

•	(TETRAHEDRAL) Structure
	Methane
Molecular Structure

• (TETRAHEDRAL) Structure
 Methane
1. Molecular Structure: The VSEPR Model

VSEPR Model

• VSEPR: Valence Shell Electron-Pair _________.
• The structure around a given atom is determined principally by __________ electron pair repulsions.
Molecular Structure: The VSEPR Model

VSEPR Model

• VSEPR: Valence Shell Electron-Pair (REPULSION).
• The structure around a given atom is determined principally by (MINIMIZING) electron pair repulsions.
2.	Molecular Structure: The VSEPR Model

Two Pairs of Electrons

•	BeCl2
•	180°
•	______
2. Molecular Structure: The VSEPR Model

Two Pairs of Electrons

• BeCl₂
• 180°
• ______
Molecular Structure: The VSEPR Model

Two Pairs of Electrons

•	BeCl2
•	180°
•	(LINEAR)
Molecular Structure: The VSEPR Model

Two Pairs of Electrons

• BeCl₂
• 180°
• (LINEAR)
3.	Molecular Structure: The VSEPR Model

Three Pairs of Electrons

•	BF3
•	120°
•	________ ______
3. Molecular Structure: The VSEPR Model

Three Pairs of Electrons

• BF₃
• 120°
• ________ ______
Molecular Structure: The VSEPR Model

Three Pairs of Electrons

•	BF3
•	120°
•	(TRIGONAL PLANAR)
Molecular Structure: The VSEPR Model

Three Pairs of Electrons

• BF₃
• 120°
• (TRIGONAL PLANAR)
4.	Molecular Structure: The VSEPR Model

Four Pairs of Electrons

•	CH4
•	109.5°
•	___________
4. Molecular Structure: The VSEPR Model

Four Pairs of Electrons

• CH₄
• 109.5°
• ___________
Molecular Structure: The VSEPR Model

Four Pairs of Electrons

•	CH4
•	109.5°
•	(TETRAHEDRAL)
Molecular Structure: The VSEPR Model

Four Pairs of Electrons

• CH₄
• 109.5°
• (TETRAHEDRAL)
5. Molecular Structure: The VSEPR Model

Steps for Predicting Molecular Structure Using the VSEPR Model

1. Draw the Lewis structure for the molecule.
2. Count the electron pairs and arrange them in the way that minimizes repulsion (put the pairs as ___ _____ as possible).
3. Determine the _________ of the atoms from the way electron pairs are shared (how electrons are shared between the central atom and surrounding atoms).
4. Determine the ____ of the molecular structure from positions of the atoms.
Molecular Structure: The VSEPR Model

Steps for Predicting Molecular Structure Using the VSEPR Model

1. Draw the Lewis structure for the molecule.
2. Count the electron pairs and arrange them in the way that minimizes repulsion (put the pairs as [FAR APART] as possible).
3. Determine the (POSITIONS) of the atoms from the way electron pairs are shared (how electrons are shared between the central atom and surrounding atoms).
4. Determine the (NAME) of the molecular structure from positions of the atoms.
6.	Molecular Structure: The VSEPR Model

Arrangements of Electron Pairs and the Resulting _________ __________ for Two, Three, and Four Electron Pairs
6. Molecular Structure: The VSEPR Model

Arrangements of Electron Pairs and the Resulting _________ __________ for Two, Three, and Four Electron Pairs
Molecular Structure: The VSEPR Model

Arrangements of Electron Pairs and the Resulting (MOLECULAR STRUCTURES) for Two, Three, and Four Electron Pairs
Molecular Structure: The VSEPR Model

Arrangements of Electron Pairs and the Resulting (MOLECULAR STRUCTURES) for Two, Three, and Four Electron Pairs
7.	Molecular Structure: Molecules with Double Bonds

Molecules with Double Bonds

•	When using the VSEPR model to predict the molecular geometry of a molecule, a ______ or ______ bond is counted the same as a single electron pair.
	CO2
7. Molecular Structure: Molecules with Double Bonds

Molecules with Double Bonds

• When using the VSEPR model to predict the molecular geometry of a molecule, a ______ or ______ bond is counted the same as a single electron pair.
 CO₂
Molecular Structure: Molecules with Double Bonds

Molecules with Double Bonds

•	When using the VSEPR model to predict the molecular geometry of a molecule, a (DOUBLE) or (TRIPLE) bond is counted the same as a single electron pair.
	CO2
Molecular Structure: Molecules with Double Bonds

Molecules with Double Bonds

• When using the VSEPR model to predict the molecular geometry of a molecule, a (DOUBLE) or (TRIPLE) bond is counted the same as a single electron pair.
 CO₂
8. Molecular Structure: Molecules with Double Bonds

Concept Check

Determine the molecular structure for each of the following molecules, and include bond angles:

HCN
PH₃
SeO₂
O₃
Molecular Structure: Molecules with Double Bonds

Concept Check

Determine the molecular structure for each of the following molecules, and include bond angles:

HCN – linear, 180°
PH₃ – trigonal pyramid, 109.5° (107°)
SeO₂ – bent, 120°
O₃ – bent, 120°
1. Most chemical bonds consist of electrostatic attractive forces and are called ____ bonds, or of shared electrons and are called _____bonds.

1. electric; shared
2. ionic; covalent
3. ionic; molecular
4. electronic; coordinate
Choice #2 properly identifies the bonds as ionic and covalent.

Section 12.1: Types of Chemical Bonds
2. When electrons in a covalent bond are shared equally the bond is _________, but if the electrons are not shared equally the bond is ________, which means that it has a positive side and a negative side.

1. ionic; covalent
2. strong; weak
3. nonpolar; polar
4. stable; unstable
Choice #3 is the correct answer.

Section 12.1: Types of Chemical Bonds
3. When considering a bond between two atoms, the greater the difference in ____________, the more __________is the bond.

1. polarity; divided
2. atomic weight; nonpolar
3. electronegativity; polar
4. electronegativity; nonpolar
Choice #3 correctly expresses the idea that bonds become more polar as the difference in the electronegativity of the atoms increases.

Section 12.2: Electronegativity
4. Rank the following bonds from least polar to most polar:

Si-Cl, P-Cl, Mg-Cl, S-Cl

1. P-Cl, S-Cl, Si-Cl, Mg-Cl
2. Mg-Cl, Si-Cl, P-Cl, S-Cl
3. Mg-Cl, S-Cl, P-Cl, Si-Cl
4. S-Cl, P-Cl, Si-Cl, Mg-Cl
Choice #3 is the correct answer. Bonds become more polar as the difference in the electronegativity of the atoms increases.

Section 12.2: Electronegativity
5. Choose the bond that is the most polar.

1. Sr–O
2. C–O
3. N–N
4. Fe–P
Choice #1 is the correct answer. Bonds become more polar as the difference in the electronegativity of the atoms increases.

Section 12.2: Electronegativity
6. Which of the following bonds would be the least polar yet still be considered polar covalent?

1. Mg–O
2. N–O
3. Si–O
4. O–O
Choice #2 is the correct answer. To be considered polar covalent, unequal sharing of electrons must still occur. Choose the bond with the least difference in electronegativity yet there is still some unequal sharing of electrons.

Section 12.2: Electronegativity
7. Which of the following bonds would be the most polar without being considered ionic?

1. Mg–O
2. N–O
3. Si–O
4. O–O
Choice #3 is the correct answer. To not be considered ionic, generally the bond needs to be between two nonmetals. The most polar bond between the nonmetals occurs with the bond that has the greatest difference in electronegativity.

Section 12.2: Electronegativity
8. The difference in electronegativity between hydrogen and iodine is (4.0 – 2.5) = 1.5, which means that hydroiodic acid has a(n) ________________.

1. low boiling point
2. low acidity
3. unstable gas phase
4. dipole moment
Choice #4 associates the bond polarity of a diatomic molecule with having a dipole moment.

Section 12.3: Bond Polarity and Dipole Moments
9. If we consider the electron configuration of strontium, [Kr]5s², and that of oxygen, [He]2s²2p⁴, both atoms will attain stable noble gas electron configurations by the transfer of ____ electron(s). This will give Sr a charge of ___ and O a charge of ___. Hence the ionic compound formed has the formula ___________ and is named strontium oxide.

1. 1 ; 1+ ; 1– ; SrO
2. 2 ; 2+ ; 2– ; SrO
3. 2 ; 2+ ; 1– ; SrO
4. 2 ; 2– ; 2+ ; SrO
Choice #2 correctly indicates the transfer of two electrons from strontium to oxygen to form the ionic compound known as strontium oxide.

Section 12.4: Stable Electron Configurations and Charges on Ions
10. What is the expected ground state electron configuration for Te²⁻?

1. [Kr]5s²4d¹⁰5p⁴
2. [Kr]5s²4d¹⁰4f¹⁴5p⁶
3. [Kr]5s²4d¹⁰5p⁶
4. [Ar]5s²4d¹⁰5p²
Choice #3 is the correct answer. Te²⁻ contains 54 electrons and has the noble gas configuration of Xe.

Section 12.4: Stable Electron Configurations and Charges on Ions
11. What is the correct electron configuration for the most stable form of the sulfur ion in an ionic compound?

1. 1s²2s²2p⁶3s²
2. 1s²2s²2p⁶3s²3p²
3. 1s²2s²2p⁶3s²3p⁴
4. 1s²2s²2p⁶3s²3p⁶
Choice #4 is the correct answer. S²⁻ is the most stable form of the sulfur ion in an ionic compound. It contains 18 electrons and has the noble gas configuration of Ar.

Section 12.4: Stable Electron Configurations and Charges on Ions
12. The structures of ionic compounds are usually described as the packing of ____________ with smaller ____________ fitting into the interstices.

1. anions; cations
2. anions; electrons
3. cations; anions
4. cations; electrons
Choice #1 correctly describes the packing of anions (which tend to be larger) into specific patterns known as crystal lattices, while cations fit into the spaces or interstices between the packed anions.

Section 12.5: Ionic Bonding and Structures of Ionic Compounds
13. Rank the following from smallest to largest atomic radius:

Ar, S²⁻, Ca²⁺, K⁺, Cl⁻

1. Ar < K⁺ < Ca²⁺ < S²⁻ < Cl⁻
2. Ca²⁺ < K⁺ < Ar < Cl⁻ < S2²⁻
3. Ar < Cl⁻ < S²⁻ < Ca²⁺ < K⁺
4. S²⁻ < Cl⁻ < Ar < K⁺ < Ca²⁺
Choice #2 is the correct answer. They all have the same electron configuration of the noble gas Ar. Therefore, nuclear charge becomes very important in determining the sizes relative to each other. The higher the nuclear charge, the smaller the ion/atom.

Section 12.5: Ionic Bonding and Structures of Ionic Compounds
14. Which atom or ion has the smallest radius?

1. O²⁺
2. O⁺
3. O
4. O²⁻
Choice #1 is the correct answer. As electrons are removed, the effective nuclear charge has a stronger effect and thus makes the ion, O²⁺, the smallest.

Section 12.5: Ionic Bonding and Structures of Ionic Compounds
15. Which of the following has the lowest ionization energy?

1. Se²⁻
2. Br⁻
3. Sr²⁺
4. Rb⁺
Choice #1 is the correct answer. Se²⁻ has the smallest effective nuclear charge and thus does not bind the electrons as strongly as the others.

Section 12.5: Ionic Bonding and Structures of Ionic Compounds
16. Lewis structures show the arrangement of ________electrons in an atom or ion.

1. all
2. core
3. valence
4. missing
Choice #3 correctly indicates that only valence electrons are included in Lewis structures.

Section 12.6: Lewis Structures
17. In the Lewis structure for H₂S there are a total of ______ electrons and ____ pair(s) of nonbonding electrons.

1. 9 ; 2
2. 9 ; 1
3. 8 ; 2
4. 8 ; 1
Choice #3 correctly describes the Lewis structure for hydrosulfuric acid: (6 electrons from S) + (2 × 1 = 2 electrons from H) = 8 total electrons. The Lewis structure is the same as that for H₂O, with S in place of O.

Section 12.6: Lewis Structures
18. The Lewis structure for SO₂ contains ____ total electrons and____ nonbonding pairs of electrons, as well as one ______ bond and one _______ bond between the central sulfur and the oxygen atoms.

1. 16 ; 4 ; single ; double
2. 16 ; 6 ; single ; double
3. 18 ; 4 ; single ; triple
4. 18 ; 6 ; single ; double
Choice #4 correctly describes the SO₂ molecule with (3 × 6 =) 18 electrons and S located between the two O’s with one single and one double bond.

Section 12.7: Lewis Structures of Molecules with Multiple Bonds
19. Consider the following compounds:

CO₂, N₂, CCl₄

Which compound exhibits resonance?

1. CO₂
2. N₂
3. CCl₄
4. At least two of the above compounds exhibit resonance.
Choice #1 is the correct answer. Only CO₂ exhibits resonance.

Section 12.7: Lewis Structures of Molecules with Multiple Bonds
20. Which of the following supports why Lewis structures are not a completely accurate way to draw molecules?

1. We cannot say for certain where an electron is located yet when drawing Lewis structures, we assume the electrons are right where we place them.
2. When adding up the number of valence electrons for a molecule, it is possible to get an odd number which would make it impossible to satisfy the octet rule for all atoms.
3. Both statements 1 and 2 above support why Lewis structures are not a completely accurate way to draw molecules.
4. Lewis structures are the most accurate way to draw molecules and are completely correct.
Choice #3 is the correct answer.

Section 12.7: Lewis Structures of Molecules with Multiple Bonds
21. It is important to fully understand that Lewis structures are useful in determining the bonding relationships between atoms in a molecule, but that they do not directly provide a true picture of molecular shape. While the Lewis structure for methane, CH4, an important greenhouse gas, suggests a flat structure with 4 hydrogens arranged around a central carbon, the methane molecule is actually ________.

1. square planar
2. trigonal
3. tetrahedral
4. octahedral
Choice #3 correctly describes the methane molecule as a tetrahedron, with the carbon centrally located in the geometric center and the 4 hydrogens forming the corners of the tetrahedron.

Section 12.8: Molecular Structure
22. While the electron pair geometry of NH3 is _______, VSEPR predicts the molecular shape as _________, due to the pair of nonbonding electrons on the central N.

1. tetrahedral; trigonal pyramidal
2. trigonal planar; tetrahedral
3. tetrahedral; trigonal planar
4. Both are trigonal planar.
Choice #1 answers both geometry questions correctly. The central N has four pairs of electrons around it, giving it a tetrahedral electron pair geometry. But because one of the corners of the tetrahedron is an electron pair, the molecule is a trigonal pyramid with the N forming the apex and the three H’s forming the pyramidal base.

Section 12.9: Molecular Structure: The VSEPR Model
23. Draw the Lewis structures for the following compounds:

CBr₂H₂, BH₃, HCl

Which compound has bond angles of 109.5˚ around the central atom?

1. BH₃
2. CBr₂H₂
3. HCl
4. At least two of the above compounds have bond angles of 109.5˚.
Choice #2 is the correct answer. CBr₂H₂ is tetrahedral, BH₃ is trigonal planar, and HCl is linear.

Section 12.9: Molecular Structure: The VSEPR Model
24. In determining the shape of the SO₂ molecule we examine the Lewis structure and find the central S atom attached to O’s via one single and one double bond. The electron pair geometry is _______ and the molecular geometry is _________, with bond angles of ____ degrees.

1. trigonal planar ; linear ; 180
2. trigonal planar ; bent ; 120
3. tetrahedral ; linear ; 120
4. tetrahedral ; trigonal planar ; 120
Choice #2 correctly describes the electron pair about the central S as trigonal planar. S has one nonbonding electron pair, a single bond, and a double bond (counts as a single bond for shape). Because one corner of the triangle is occupied by an electron pair, the molecule has a bent shape with bond angles of 120˚.

Section 12.10: Molecular Structure: Molecules with Double Bonds
25. Draw the Lewis structures for the following compounds:

HCN, NH₄⁺, NO₂⁻

Which compound has bond angles of 120˚ around the central atom?

1. HCN
2. NH₄⁺
3. NO₂⁻
4. At least two of the above compounds have bond angles of 120˚.
Choice #3 is the correct answer. NO₂⁻ has trigonal planar electron pair geometry. NH₄⁺ has tetrahedral electron pair geometry and HCN has linear electron pair geometry.

Section 12.10: Molecular Structure: Molecules with Double Bonds
26. How many of the following statements are true regarding the molecule SeO₂?

I. The electron pair geometry of the molecule is trigonal planar.
II. The molecular structure of the molecule is linear.
III. The number of effective pairs around the central atom is four.

1. 0
2. 1
3. 2
4. 3
Choice #2 is the correct answer. Only statement I is true. SeO₂ has trigonal planar electron pair geometry, bent molecular structure, and three effective pairs around the central atom S.

Section 12.10: Molecular Structure: Molecules with Double Bonds