• 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

1/98

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;

98 Cards in this Set

  • Front
  • Back
CHAPTER ONE

What is Chemistry?
The science that deals with the materials of the universe and the changes that these materials undergo. Examples: wood burning and forming water, carbon dioxide and other substances. Steel rusting; rocket-ship fuel reacting to propel the space shuttle.
Problem-solving process for everyday problems:
1. Recognize the problem and state it clearly (make an observation).
2. Propose possible explanations to the problem or possible explanations for the observation (formulate a hypothesis).
3. Decide which is the best solution or the most reasonable explanation (perform an experiement to obtain new information).
Scientific Method for solving problems:
1. State the problem and collect data (make observations).
-Qualitative observation: does not involve a number. Ex: the sky is blue.
-Quantitative observation: involves a number. Ex: water boils at 100 degrees C.
2. Formulate Hypothesis: a hypothesis is a possible explanation for the observation.
3. Perform experiments: experiments are something we do to produce new information to test the hypothesis.
What is a theory (model)? Observations?
Theory (model)-A set of tested hypotheses that gives an overall explanation of some part of nature.

Observation is something that is witnessed and can be recorded.
-Observations do not change, but the theories (interpretations and human inventions) many change as more information becomes available.
Natural Law? Law of Conservation of Mass?
Natural law is a generally observed behavior applying to many different systems.
The Law of Conservation of Mass is the observation that the total mass of materials in not affected by a chemical change.
The difference between a law and a theory?
The difference between a law and a theory is that a law is a summary of observed (measurable) behavior; whereas, a theory is an explanation of behavior.
The scientific method is only as effective as the humans using it and does not automatically lead to progress.
CHAPTER TWO

Two types of observations, quantitative and qualitative.
Qualitative observation: Measurements consisting of a number and a unit.
Qualitative observation: Observations that do not contain a number. Ex-the powder is white.
Scientific notation: Scientific notation is a method for making very large or very small numbers more compact and easier to write.
Scientific notation expresses a number as a product of a number between 1 and 10.

IF THE DECIMAL POINT IS MOVED TO THE LEFT, THE EXPONENT OF 10 IS POSITIVE. Ex: 0.000167 in scientific notation is 1.67 x 10 to the fourth power.

Ex: 93,000,000 = 9.3 x 10,000,000 = 9.3 x 10 to the 7th power (a number between 1 and 10, times a power of 10).
Using scientific notation:
-Any number can be represented as the product of a number between 1 and 10 and a power of 10.
-The power of 10 depends on the number of places the decimal point is moved and in which direction. The number of places the decimal point is moved determines the power of ten. The direction of the move determines whether the power of 10 is positive (moved to the left) or negative (moved to the right).
Units and systems of units; English used in the US or Metric used in most of the industrialized world and in science.
A unit is the part of a measurement that tells what scale or standard is being used to represent the results of the measurement.

In 1960 and international agreement set up a comprehensive system of units called the International System (SI) based on the metric system.
Mass-kilogram-kg
Length-meter-m
Time-second-s
Temperature-Kelvin-K
Kilo-k-1000-10 to the 3rd
Deci-d-0.1-10 to the -1
Centi-c-0.01-10 to the -2
Milli-m-0.001-10 to the -3
Micro-u-0.000001-10 to the -6
Nano-n-0.000000001-10 to the -9
Measurements of Length:
Fundamental SI unit of length is a meter.
-one meter is 39.37 inches and 2.54 cm = one inch.
-prefixes can be used to express fractions of multiples of a meter. Ex: 1 kilometer = 1000 meters (1km) or 1 centimeter (1cm) = 0.01 meters (m).
Measurements of Volume:
Fundamental SI unit of volume is a liter.
-volume is the amount of three-dimensional space occupied by a substance.
-one liter = 1L=1000 ml and 1 milliliter = 1mL = 10 to the -3 power L. Also, 1 milliliter (mL) = 1 cm to the 3rd.
-Graduated Cylinder - used to measure volume in the laboratory.
Meaurements of Mass:
Fundamental SI unit of mass = kilogram.
-a balance used to determine the mass of an object.
-a kilogram = 1 kg = 1000g = 10 to the 3rd g.
-1 gram = 1g (Fundamental unit in metric system).
-1 milligram = 1 mg = 0.001 g = 10 to the negative 3 g.
Uncertainty in measurements:
Whenever a measurement is made with a device, like a ruler, volumetric, etc., an estimate is required and the measurement has some degree of uncertainty.
Ex: Measuring a pin with a ruler with 0.1 cm markings.
-the pin measure between 2.8-2.9 cm.
-we estimate that the pin in 2.85 cm long, but it could be 2.84 or 2.86 cm.
Significant Figures: The numbers recorded in a measurement (all the certain numbers plus the first uncertain number).
The number of significant figures for a given measurement depends on the inherent uncertainty of device used to take the measurement.
-The uncertainty in the last number (estimated number) is usually assumed to be +/= 1, unless otherwise indicated.
-The length of the pin measured above could be written 2.85 +/= 0.01 cm.
SIGNIFICANT FIGURES: Chemistry involves using measured numbers in calculations and it is important to know what happens when we do math with numbers that contain uncertainties.

Rules for Counting Significant Figures:
1. Nonzero integers always count as significant figures (between 1 & 9). Ex: 1,457 contains 4 significant figures.
2. Zeros: there are three classes of zeros.
a. Leading zeros are zeros that precede all of the nonzero digits. These numbers never count as significant figures. Ex: 0.0025 contains only two significant figures (the three zeros only indicate the position of the decimal point).
b. Captive zeros are zeros that fall between nonzero digits. These numbers always count as significant figures. Ex: 1.008 has four significant figures.
c. Trailing zeros are zeros at the right end of the number. They are significant only if the number is written with a decimal point. Ex: 100 has only 1 SF and 100. has 3 SF's.
More examples of counting SF's:
-0.0108g (3 SF's)
-0.0050060 (5 SF's)
-5.030 x 10 to the 3rd ft. (4 SF's)
-50 students (exact number and an unlimited number of significant figures).
-204,150 (5 SF's)
-1430 (3 SF's) written as 1.43 x 10 to the 3rd.
-1430. (4 SF's) written as 1.43 x 10 to the 4th.
****if you put a decimal point in the measurement, it implies a number with error figured in. 1000.g means that there are 999-1001 grams.
**If there is a decimal point in a number, count trailing zeros. If there is no decimal point, DO NOT count trailing zeros.
Rounding off numbers:
Your calculator usually displays too many digits and the number needs to be "rounded off" to the correct number of significant figures.
Rules for rounding off:
1. If the digit to be removed
a. Is less than 5, the preceding digit stays the same. Ex: 1.33 rounds to 1.3.
b. Is equal to or greater than 5, the preceeding digit is increased by 1. Ex: 1.36 rounds to 1.4 and 3.15 rounds to 3.2.
2. In a series of calculations, carry the extra digits through to the final result and then round off.
**When rounding off, use onlly the first number to the right of the last SF to determine which direction to round off.
-Ex: 4.348 is rounded off to 4.3 when there are 2 SF's. 4.348 is rounded to 4.35 when there are 3 SF's.
Significant Figures in Calculations:
1. For multiplication and division, the number of SF's in the result is the same as that in the measurement with the smallest number of SF's (limiting measurement). Ex: 4.56 (2 SF's) x 1.4 (2 SF's) = 6.383. Round off to 6.38 (3 SF's) = 6.4 (2 SF's). Or, 8.315 / 298 = 0.0279027. Round off to 2.79 x 10 to the -2.
2. For addition and subtraction, the limiting term is the one with the smallest number of decimal places. Ex: ADDITION: 12.11 + 18.0 + 1.013 = 31.123, round to 31.1. SUBTRACT: 0.6875 - 0.1 = 0.5875, round off to 0.6
Conversion Factors:
*One unit can be converted to another unit by means of a conversion factor. Unit(1) x conversion factor = Unit(2).
*An Equivilance Statement relates the same quantity in two different units. Ex: 2.54 cm = 1 in and 1 mi = 5280 feet.
*Conversion Factor is a ratio of parts of the statement that relates the two units or a ratio of the two parts of the equivilance statement. Ex: 2.54 cm / 1 in or 1 in / 2.54 cm.
English-Metric and English-English conversions:
Length
-1 m = 1.094 yd
-2.54 cm = 1 in.
-1 mi = 5280. ft
-1 mi = 1760. yd
Mass
-1 kg = 2.205 lb
-453.6 g = 1 lb (rounded to 454 g = 1 lb)
Volume
-1 L = 1.06 qt

Convert 2.85 cm into inches:
2.85 cm / 1 x 1 in / 2.54 cm = 1.12 in (3 SF's because 2.85 is limiting).
Convert 7.00 in to cm's. 7.00 in / 1 x 2.54 cm / 1 in = (7.00)(2.54) cm = 17.8 cm (3 SF's because 7.00 is limiting).
Dimensional Analysis:
Dimensional Analysis is changing from one unit to another via conversion facts (based on the equivilance statments between the units).
*Remember to move decimal to make 42g x 10 to the 3rd mg / 1g = 4.2 x 10 to the 4th. 1 g = 10 to the 3rd mg.
Converting from one unit to another:
Step 1: To convert from one unit to another, use the equivilance statement relating to two units. The conversion factor is a ratio of the two parts of the equivilance statement.
Step 2: Choose the appropriate conversion factor by looking at the direction of the required change (make sure the unwanted units cancel).
Step 3: Multiply the quantity to be converted by the conversion factor to give the quantity with the desired units.
Step 4: Check that you have the correct number of significant figures.
Step 5: Ask whether your answer makes sense.
Conversion Factors: Multiple Step Problems
Ex: The length of the the marathon race is approximately 26.2 miles. What is the distance in km?

Miles-years-meters-kilometers.

Miles-yards: 26.2 mi x 1760 yd / 1 mi = 46,112 yd = 46,100 = 4.61 x 10 to the 4th yd (3 SF's).

Yards-meters: 4.61 x 10 to the 4th yds x 1 km / 10 to the 4th m = 42.1 x 10 to the 4th yds (3 SF's).

Meters-kilometers: 4.21 x 10 to the 4th m x 1 km / 10 to the 3rd m = 4.21 x 10 to the first km = 42.1 km (3 SF's).
Temperature Conversions:
Fahrenheit Scale: Used widely in the US and Great Britain and in more engineering sciences. Ex: The outside temperature is 75 degrees and water freezes at 32 degrees and boils at 212 degrees F.
Celcius Scale: Used in Canada and Europe and in the physical and life sciences in most countries. Ex: In keeping with the metric system, the freezing and boiling points of water are 0 degrees and 100 degrees C.
Absolute or Kelvin Scale: Used in the sciences. Ex: Water freezes at 273 degrees and boils at 373 degrees K.
Comparing the three temperature scales
1. The size of each temperature unit (each degree) is the same for the celcius and kelvin scales, since the difference between the boiling and freezing points of water is 100 units on both of these scales.
2. The F degree is smaller than the C and K units. On the F scales there are 180 degrees between boiling and freezing points of water, as compared with 100 units on the other two scales.
3. The zero points are different on all three scales.
Converting between temperature scales:
Celcius to Kelvin:
Temp in kelvin = temp in celcius + 273.

Kelvin to celcius:
Temp in C = Temp in K - 273.

Celcius to fahrenheit:
Temp in degrees F = 1.80(temp in C) + 32

Fahrenheit to Celcius:
Temp in C = temp in F - 32 / 1.80
Conversions, summarized:
C to K: K = C + 273
K to C: C = K - 273
C to F: F = 1.80(C) + 32
F to C: C = F - 32 / 1.80
Density
Density = mass / volume

*the density of a liquid is determined by weighing a known volume of hte liquid. Ex: if 23.50 mL of a certain liquid weighs 35.062 g, than what is it's density? 35.062 / 23.50 mL = 1.492 g/mL or 1.492 g/cm to the 3rd. One cm to the third is equal to 1 mL.
*The volume of an irregular solid object is often determined indirectly by submerging it in water and measuring the volume of the water displaced.
Using density in calculations
Mercury has a density of 13.6 g/mL. What volume of mercury must be taken to obtain 225 g of the metal? v=m/d; 225g / 3.6 g/mL = 16.5 mL. d/1 = m/v, cross multiply, d x v = m x 1, v = m / d.

*Density can be used to determine whether your car battery needs to be recharged (sulfuric acid and water have different densities) and whether you need more antifreeze (antifreeze and water have different densities).

HYDROMETER is used to determine the density of liquids.
CHAPTER 3: Matter
Matter is the stuff that composes the universe and has two characteristics - it has mass and it occupies space.

3 STATES OF MATTER: SOLID, LIQUID AND GAS

Solids are rigid; they have a fixed shape and volume. Ex: ice cube, diamond, iron bar.

Liquids have a definite volume, but take the shape of the container. Ex: gasoline, water, alcohol and blood.

Gas has no fixed volume or shape; it takes the shape and volume of it's container. Ex: air, helium and oxygen.

ENERGY DOES NOT HAVE MATTER AND DOES NOT OCCUPY SPACE.
Physical Properties and their changes:
Physical properties have characteristics identifying a person (e.g. hair color, eye color).
Physical properties of substances are physical characteristics that identify the substance (e.g. color, odor, state (sold, liquid or gas), density, boiling point, etc.).
Chemical Properties & physical property changes:
Chemical properties are a substances ability to form new substances, like wood burning, steel rusting, grass growing, etc. Examples:

a. the boiling point of a certain alcohol is 78 C-physical property.
b. diamond is very hard-physical property.
c. sugar ferments to form alcohol-chemical property (sugar chamically changes into alcohol).
d. a metal wire conducts an electric current-physical property.

Qualitative observations:

e. Gallium metal melts in your hand-physical property.
f. Platinum does not react with oxygen at room temperature-chemical property.
g. The paper is white-physical.
h. The copper sheets on the statue of liberty have acquired a greenish coating over the years-chemical.
WATER:
Water molecules are made up of two atoms of hyrogen and one atom of oxygen (H20 or H-O-H).

Physical changes do not affect the composition of the substance. Water goes from ice (solid) to water (liquid) to steam (vapor).

Chemical changes-A substance forms a new substance. The electrolysis of water forms a hydrogen molecules and oxygen molecules.
More re: physical, chemical changes and reactions.
Physical changes involves a change in one or more physical properties, but no change in the fundamental components that make up the substance. The most common physical changes are changes of state: solid-liquid-gas.
a. solid state is the when the molecules are close together and are locked into rigid positions.
b. liquid state is when the molecules are still close together, but can move around.
c. gaseous state is when the molecules are far apart and move randomly.

Chemical change involves a change in the fundamental components of the substance; a given substance changes into a different substance or substances. Ex: passing an electric current through water changes it in to hydrogen gas and oxygen gas. (2)H20 = (2)H2 + 02.

Reactions are chemical changes. Ex: silver tarnishes by reacting with substances in the air.
Elements and Compounds
Element: A substance that cannot be broken down into the other substances by chemical methods. Ex: iron, aluminum, oxygen, and hydrogen. AN ELEMENT IS COMPOSED OF ATOMS OF THE ELEMENT ONLY.

Compound: A substance composed of a given combination of elements that can be broken down into those elements by chemical methods: Compounds-Elements. -Ex: water is a compound consisting of H2) molecules. COMPOUNDS ALWAYS CONTAIN ATOMS OF A DIFFERENT ELEMENTS; HOWEVER, COMPOUNDS ALWAYS HAVE THE SAME COMPOSITION.

-The properties of a compound are usually different from those of the elements it contains. Hyrdrogen + Oxygen (gases) form water(liquid). NACL: table salt = (NaCl).
Mixtures
Mixtures are things that have variable composition.
-almost all the matter around us are mixtures of substances.
Ex: Wood-composition varies with the tree. Wine-compostion varies with the grapes used-white, red, etc. Coffee-composition varies with the types of coffee used and the amounts. Well water-composition varies with the dissolved minerals and gases.

Distilled water would be a pure substance.

Mixtures can be separated into pure substances: elements and/or compounds.

Mixtures had 2 or more pure substances.

Pure Substancesare elements or compounds that always have the same composition. Ex: pure water, pure copper or pure aluminum.
Two type of mixtures: Homo and heterogeneous.
Homogeneous mixture is the same throughout and is also called a solution. Ex: salt dissolved in water, brass, air are homogeneous mixtures or solutions.

Heterogeneous mixtures contain regions that have different properties from those of other regions. Ex: sand in water, salt and pepper, oil and water, salt & sugar.

gasoline-homo
chocolate chip cookies-hetero
oil & water salad dressing-hetero
brass-homo
air-homo
common salt (sodium chloride) compound-pure
pure apple juice-homo
oxygen gas (element)-pure
Methods of separating mixtures
Distillation-purifying a liquid from a dissolved solid by boiling the liquid, collecting the vapor and condesing it. The liquid is purified and the solds are left behind. Ex: Purifying salt water. The water changes into a gas, which travels through a condensing tube and is cooled and turns again into a liquid. The sodium chloride remains unchanged and remains behind.

Filtration is the separation of a liquid from a solid by passing the liquid through filter paper, but the solid particles are left behind on the paper. Ex: separating sand from water. The sand remains behind the filter paper.
The organization of matter:
MATTER:

HOMOGENEOUS OR HETEROGENEOUS MIXTURE, WHICH ARE PHYSICAL METHODS.

PURE SUBSTANCES ARE EITHER ELEMENTS OR COMPOUNDS AND ARE CHEMICAL METHODS.
CHAPTER 4:

Chemical Foundations: Elements, Atoms and Ions.
*Man has used chemical processes, like processing ores to produce metals for making tools, etc. before 1000 BC
*By 400 BC, the Greeks proposed that all matter was composed of four fundamental substances; fire, earth water and air.
*Alchemy dominated the next 2000 years with some fakes and some sincere scientists.
*Robert Boyle (1627-1691) was the first scientist to recognize the importance of careful measurements and conducting experiemnts and is known for his pioneering work on gases. He defined an element as a substance that could not be broken down into many pure substances.
The Elements:
There are 115 different elements known today, with 88 of them occurring naturally, and millions of known substances are made from the elements.

*The following distribution (mass %) of the most abundant elements in the earth's crust, oceans and atmosphere.

Oxygen, 49.2%, Silicon, 25.7% (sand is silicon dioxide), Aluminum, 7.5%.

*Oxygen accounts for 20% of the earths atmosphere (O2) and is found in virtually all the rocks, sand and soil on the earth's crust.
*The distribution of elements found in living matter is very different from the list above.

Oxygen, 65%, Carbon, 18%, Hydrogen, 10%. (Organic chemicals are cabon and hydrogen).
Trace elements & elements:
Trace elements are elements such as chromium, cobalt, iodine, manganese and copper. They are all necessary for life, even though they are present in only small amounts.

Element is a term that is used in many ways.
1. A single atom of the at element-microscoptic form.
2. A sample of the element large enough to weight which contains many atoms of the element-macroscropic form. Elements many contain molecules instead of atoms (O2 or N2).
3. Generically, to mean the element is present in some form which may not be the free form.
Symbols for the Elements
The names for the chemical elements have been derived from the different soures.
-A Greek, Latin, or German word that describes some property of the element (i.e. gold=Au from aurum, a Latin word meaning shining dawn).
-The place where it was discovered (i.e. francium, germanium, californium).
-Scientists (i.e. einsteinium and nobelium).
Element Symbols
A set of abbreviations for the elements usually consisting of the first letter or the first two letters of the element name. Only the first letter is capitalized:

Fluorine-F, Oxygen-O, Neon-Ne, Silicon-Si.

Some elements symbols are no the first two letters in the name:

Zinc-Zn, Chlorine-Cl, Cadmium-Cd, Platinum-Pt, are examples.

Some elements have symbols based on the original Latin or Greek name:

Gold-aurum-Au, Lead-plubum-Pb, Sodium-natrium-Na, Iron-ferrum-Fe.
Elements to know: part one
Al - Aluminum
Ba - Barium
Br - Bromine
Ca - Calcium
C - Carbon
Cl - Chlorine
Cr - Chromium
Cu - Copper
F - Fluorine
Au - Gold
He - Helium
H - Hydrogen
I - Iodine
Fe - Iron
Pb - Lead
Observations made by scientists:
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 (the law of constant composition).
Dalton's Atomic Theory
Proposed by the English scientist, John Dalton, in 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.
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. (Water is the same in MI as it is in France).
5. Atoms are indivisible in chemical processes and are not created or destroyed in chemical reactions. A chemical reaction simply changes the way the atoms are grouped together.
Law of constant composition:
Dalton's model successfully explained the law of constant composition because if a compound contains the same relative number of atoms, it will always contain the same proportions by mass of the various elements.

Dalton's model also predicted how a given pair of elements can combine to form more than one compound. Ex: Nitrogen and oxygen can combine in different ratios to form many different compounds (NO, N2O, NO2).
Formulas of compounds:
Compounds: A distinct substance that is composed of the atoms of two or more elements and always contains exactly the same relative masses of those elements, or the same relative numbers of atoms of each element.

Chemical Fomula: A method of expressing the types of atoms and the number of each type in each unit (molecule) of a given compound.
Rules for Writing Formulas:
1. Each atom present is represented by it's 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 suscript 1 is not written.
Writing Formulas of Compounds:
Write the formula for each compound, listing the elements in the order given.
a. Each molecule of a compound that has been implicated in the formation of acid rain contains one atom of sulfur and three atoms of oxygen (SO 3).
b. Each molecule of a certain compound contains two atoms of nitrogen and five atoms of oxygen (N2 O5).
c. Each molecule of glucose, a type of sugar, contains six atoms of carbon, twelve atoms of hydrogen, and six atoms of oxygen (C6 H12 O6).
The Structure of the Atom
In the late 1890's, the English physicist J.J. Thompson showed that atoms of any element can be made to emit tiny negative particles (eletrons), which were repelled by a negatively-charged electric field.
-Thompson invented the cathode ray tube and discovered that atoms contain electrons.
-Thompson also proposed that atoms contained positive particles to balance out the charge, since atoms were neutral particles.
-Thompson's cathode ray tube (a gas-filled, sealed glass tube in which a stream of electrons can pass between 2 electrodes, causing a glow between the plates) is the basis for neon lights, TV, and computer screens.
William Thompson (Lord Kelvin) and the "Plum Pudding" model:
The Plum Pudding Model of the Atom was proposed by William Thompson (Lord Kelvin) somewhere between late 1890's and 1910.
-An atom consists of a uniform "pudding" of positive charge with enough negative electrons "raisins" scattered within to balance that positive charge. *This theory was disproved.*
Ernest Rutherford
In 1911, Ernest Rutherford proposed the nuclear model of an atom.
-Rutherford noticed that some alpha particles (positively-charged particles with a mass~7500 x that of an electron) were deflected by something in the air.
-Rutherford designed an experiment in which he bombarded a thin metal foil wiht alpha particles and recorded where the alpha particles passed through the foil.
-Most of the particles passed straight through, but some were deflected at large angles or reflected backward. The "plum pudding" model did not fit the data.
-Rutherford proposed that an atom had a dense, positive center (nucleus) around which tiny electrons moved in a space that was otherwise empty (a nuclear atom).
-By 1919, Rutherford concluded that the nucleus contained protons having the same magnitude of charge as the electron, but positive instead.
-In 1932, Rutherford and James Chadwick proposed that most nuclei also contain neutral particles, called neutrons, which are only slightly more massive than a proton, and has no charge.
Modern Nuclear Atom
The modern nuclear atom:
-Consists of a tiny nucleus (~diameter 10 the -13th cm) and electrons move about or orbit around the nucleus at a distance of ~ 10 to the -8th cm.
-Consists of positively-charged protons and neutral neutrons in the nucleus.
-All atoms contain the same components, but have different chemical properties due to the number and arrangement of the electrons.
-The electrons move around in a large, mostly empty space, around the nucleus and can intermingle when atoms combine to form the molecules.
The Mass and Charge of the Electron, Proton and Neutron
Particle: Electron, Proton, Neutron.
Relative Mass: 1, 1836, 1839.
Relative Charge: 1-, 1+, 0.

*The electron is arbitrarily assigned a mass of 1 for comparison.
Isotopes
In the early 1900's, James Chadwick discovered that the nuclei of most atoms contain neutrons as well as protons, and the Dalton model of the atom was modified to "All atoms of the same element contain the same number of protons and electrons."

Isotopes: Atoms with the same number of protons, but different neutrons.

Atomic Number: The number of protons in a nucleus (& electrons).

Mass Number: The sum of the number of neutrons and the number of protons in a given nucleus.
Symbol for an Atom
A
X
Z : where X = element symbol, A=mass number, and Z=atomic number.

Ex: sodium, Na
23
Na
11 : Sodium has an atomic number of 11 and a mass number of 23.

*mass number less atomic number = number of neutrons.
Interpreting the symbol for isotopes:
In nature, elements are usually found as a mixture of isotopes. Three isotopes of elemental carbon are carbon -12, carbon -13 and carbon -14. Determine the number of protons and neutrons in the nucleus of each type of carbon atom.

Solution: The number of protons and electrons is the same for all three isotopes and is given by the atomic number of carbon, 6.

Number of Neutrons = A-Z
*Carbon-12: number of neutrons=A-Z=12-6=6
*Carbon-13: number of neutrons=A-Z=13-6=7
*Carbon-14: number of neutrons=A-Z=14-6=8
Intro to the periodic table:
*Periodic table-The elements are organized in increasing atomic number and in specific horizontal rows and vertical columns.
*The Russian scientist, Dmitri Mendeleev, arranged the elements according to similarities in their chemical properties of various "families" of elements.
*In the periodic table, the symbols for the elements are written and their atomic number is written above the symbol.
*Group-Families of elements with similar chemical properties that lie in the same vertical column on the periodic table.
Groups in the periodic table:
Each group is assigned a number, followed by the letter A.
-group 1A is Alkali Metals
-group 2A is Alkaline Earth Metals
-group 7A are the Halogens
-group 8A are Noble Gases
-group 3-12 are Transition Metals

*Most elements are metals. Metals are shown to the left of the heavy "stair step" black line on the periodic table, except for hydrogen.
Physical Properties of Metals
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.

Nonmetals-Elements that lack those properties that characterize metals.

Nonmetals lie to the right-hand side of the heavy "stair-step" black line in the periodic table.

Metalloids, or Semimetals-Elements that lie close to the "stair-step" line and have a mixture of metallic and nonmetallic properties.
Natural States of the Elements
*Most matter consists of mixtures, which contain compounds in which atoms from different elements are bound together.
*Most elements are reactive and combine with other elements to form compounds.
*Noble Metals-Unreactive metals such as gold, silver, and platinum, which exist as single atoms.
*Noble gases-Group 8 elements also appear in nature in their uncombined forms, since their atoms do not combine readily with other atoms and exist as individual atoms. Ex: He, Ne, Ar, Kr, Xe, Rn (group VIIIA).
*Diatonic molecules-Molecules made up of two atoms, like H2, N2, and O2, Cl2, F2, Br2, and I2. MEMORIZE these Group 7, DIATOMIC MOLECULES.

-A molecule of oxygen is O2 and an oxygen atom is O.

*Solid Metals-Spherical atoms are packed closely together.
*Allotropes-the different forms of the same element. Allotropes have different properties. Ex: 3 allotropes of carbon are diamond, graphite and buckminsterfullerene.
Ions
-Atoms have zero net charge because they have equal numbers of protons (+charge) and electrons (-charge).
-Ions are charged species produced by taking a neutral atom and adding or removing one or more electrons.
-Cations are ion produced when when one or more electrons are lost from a neutral atom. EX: Na (11e) Na+ (10e) +1 e- (both species have 11 protons) or Al (13 e) Al3+ (10e) + 3e- (both species have 13 protons).
-Cations are named using the name of the parent atom. Examples: Na+ is called sodium ion or sodium cation; Al3+ is called aluminum ion or aluminum cation.
-Anions are Ion produced when one or more electrons are gained by a neutral atom. EX: Cl (17 e) + e- ~Cl- (18 ) both species contain 17 protons.
.
Naming Ions, Cations and Anions
-Anions are named by taking the root of the atom name plus the suffix -ide. EX: Chloride ion Cl- is chlorine, Fluoride ion F- is fluorine, Iodide ion I- is iodine and Oxide O2- is oxygen.
*Ions are always formed by removing or adding electrons to an atom and are never formed by changing the number of protons in an atom's nucleus.
*Isolated ions do not form ions on their own, but form when metallic elements combine with non metallic elements.
Ion Charges and the Periodic Table
It is important to memorize the group numbers and the types of ions formed by members of groups 1,2, 3, 6 & 7.

Group 1 has a + charge
Group 2 has a +2 charge
Group 3 has a +3 charge
Group 5 has a -3 charge
Group 6 has a -2 charge
Group 7 has a -3 charge
Compounds that contain Ions (metal reacted wiht a nonmetal):
Ionic compounds are compounds tht contain ions.
-Since their net charge is zero, there must be both positive ions (cations) and negative ones (anions).
-The numbers of cations and anions must be such that the net charge is zero.
*Total charge of cations + Total chage of anions = Zero Net Charge.

EX: NaCl contains 1 Na+ and 1 Cl-; MgCl2 contains 1 Mg2+ and 2 Cl-; Li3N contains 3 Li+ and 1 N3-; these all become neutral.

*Ionic compounds conduct an electric current in the molten and when dissolved in water.
*Metals are always listed first, nonmetals second.
CHAPTER 5:

Common Names
Common Names-Early chemists coined names like quicklime and laughing gas because there was no organized method for naming compounds. We now have more than 4 million compounds and a system was developed for naming compounds in which the name tells something about the composition of the compound.
Naming Compounds:
-System for naming binary compounds (compounds composed of two elements).

Two type of Binary Compounds:
1. Compounds that contain a metal and a nonmetal (ionic).
2. Compounds that contain two nonmetals (molecular). *these share electrons*
Naming Compounds that Contain a Metal and a Nonmetal (types I and II):
-Binary Ionic Compound-compounds containing a positive ion (cation) and a negative ion (anion).
-The cation is written first in the formula and we name the ions.
-Type I Cations are formed from metal atoms that can form only one cation (one type). Ex: Na+, Ca2+, Cs+, and Al3+.
-Type II cations are formed from metal atoms that can form two or more cations. Ex: Cr can form Cr2+ and Cr3+ cations; Cu can form Cu+ and Cu2+ cations.
-Type I compounds contain a type I cation.
-Type II compounds contains a type II cation.
Common Simple Cations and Anions: MEMORIZE
CATIONS: TYPE ONE
Hydrogen=H+
Lithium=Li
Sodium=Na+
Potassium=K+
Cesium=Cs+
Beryllium=Be2+ Magnesium=Mg2+ Calcium=Ca2+
Barium=Ba2+
Aluminum=Al3+

*transition metals*
Silver=Ag+
Zinc=Zn2+
Rules for naming TYPE I ionic compounds:
1. The cation is always named first and the anion second.
2. A simple cation (obtained from a single atom) takes its name from the name of the element. EX: Na+ is called sodium in compounds containing this ion.
3. A simple anion (obtained from a single atom) is named by taking the first part of the element name (the root) and adding -ide. Ex: Cl- ion is called chloride.
EXAMPLES OF COMPOUNDS:
NaCl = Na+, Cl- = Sodium chloride
KL = K+, I- = Potassium iodide
CaS = Ca2+, S2- = Calcium sulfide
CsBr = Cs+, Br- = Cesium bromide
MgO = Mg2+, O2- = Magnesium oxide
CsF = Cs+, F- = Cesium fluoride
AlCl3 = Al3+, Cl- = Aluminum Chloride
MgI2 = Mg2+, I- = Magnesium Iodide
Rb2O = Rb+, O2- = Rubidium Oxide
SrI2 = Sr2+, I- = Strontium iodide
K2S = K+, S2- = Potassium sulfide

*In formulas for ionic compounds, simple ions are represented by the element symbol: Na means Na+ and CL means CL-. You write them as NaCl (they are written as neutral).
*When individual ions are shown, the charge is always included.
Ionic compounds:
1. Compounds formed from metals and nonmetals are ionic.
2. In an ionic compound the cation is always named first.
3. The net charge on an ionic compound is always a zero.
TYPE II Binary Ionic Compounds:
Since Type II binary ionic compounds contain metals that can form more than one type of cation, then multiple compounds can be made from the same two elements.

EX: FeO for Fe2+ and Fe2 O3 for Fe3+ : Fe2+ = FeO (iron (II) oxide), Fe3+ = Fe2 O3 (iron (III) oxide).

A way of specifying which type of cation is present in compounds containing metals that can form more than one type of cation is needed.

**ALL OF THE TRANSITION METALS, EXCEPT ZINE AND SILVER**

-Type II cations can have different formulas.

Roman numerals are used to specify the charge of the cation. EX: FeO is called iron (II) oxide and Fe2 O3 is called iron (III) oxide.
Common Type II Cations
Fe3+ = Iron (III) = Ferric
Fe2+ = Iron (II) = Ferrous
Cu2+ = Copper (II) = Cupric
Cu+ =Copper (I) = Cuprous
Sn4+ = Tin (IV) = Stannic
Sn2+ = Tin (II) = Stannous

*The older system for naming ionic compounds containing metals that form two cations.
-The ion with the higher charge has a name ending in -ic, and the one with the lower charge has a name ending in -ous (make sure to understand the rule).
-This system will not be used in this course, but you should be aware of it.
GROUP 1 & 2 METALS ARE ALWAYS TYPE I AND TRANSITION METALS ARE ALMOST ALWAYS TYPE II:

Rules for Naming Type II Ionic Compounds:
1. The cation is always named first and the anion second.
2. Because the cation can assume more than one charge, the charge is specified by Roman numeral in parenthesis.

TYPE II COMPOUNDS:
CuCl = Cu+, Cl- = Copper (I) chloride
HgO = Hg2+, O2- = Mercury (II) oxide
Fe2 O3 = Fe3+, O2- = Iron (III) oxide
MnO2 = Mn4+, O2- = Manganese (IV) oxide
PbCl4 = Pb4+, Cl- = Lead (IV) chloride

*charge in the metal is what determines the roman numeral.
*metals that form only one cation are not identified by a roman numeral.

COMMON METALS THAT DO NOT REQUIRE ROMAN NUMERALS ARE:

-group 1 elements which form only 1+ ions.
-gorup 2 elements which form only 2+ ions.
-some group 3 metals, like aluminum and gallium, which form 3+ ions.
ANIONS AND THEIR CHARGES
THE CHARGE ON THE METAL CAN BE DETERMINED BY BALANCING THE POSITIVE AND THE NEGATIVE CHARGES OF THE COMPOUND. YOU MUST BE ABLE TO RECOGNIZE THE COMMON ANIONS AND KNOW THEIR CHARGES.

CoBr2 = type II = Co2+, cobalt (II), Br-, bromide = Cobalt (II) bromide.

CaCl2 = type I = Ca2+, calcium, Cl-, chloride = Calcium chloride

Al2O3 = type I = Al3+, aluminum, O2-, oxide = Aluminum oxide

CrCl3 = type II = Chromium (III), Cl-, chloride
Flow chart for naming binary ionic compounds:
Ask these questions:

Does the compound contain type I or type II cations?

Type I: Name the cation, using the element name.
Type II: Transition elements (except Zinc and Silver).
* Using the principle of charge balance, determine the cation charge.
*Include in the cation name, a roman numeral indicating the charge on the cation.
Naming binary compounds that contain only nonmetals (type III):
-The system for naming binary compounds that contain only nonmetals is similar in some ways to the rules for naming binary compounds.
-Type III binary compounds contain only nonmetals.
-Rules for naming type III binary compounds:
1. The first element in the formula is named first, and the full element name is used.
2. The second element is named as though it were an anion.
3. Prefixes are used to denote the numbers of atoms present.
4. The prefix mono-is never used for naming the first element. For example, CO is called carbon monoxide, not monocarbon monoxide.

*GROUP 5 FORMS -3 IONS*
Table of prefixes for type III binary compounds:
Mono- 1
Di- 2
Tri- 3
Tetra- 4
Penta- 5
Hexa- 6
Hepta- 7
Octa- 8
Naming type III binary compounds:
BF3, boron fluoride, none/tri, boron trifluoride
NO, nitrogen oxide, none/mono, nitrogen monoxide
N2O5, nitrogen oxide, di/penta, dinitrogen pentoxide
CCl4, carbon chloride, none/tetra, carbon tetrachloride
NO2, nitrogen oxide, none/di, nitrogen dioxide
IF5, iodine fluoride, none/penta, iodine pentafluoride

*to avoid awkward pronunciation, the final o or a of the prefix is often dropped when the second element is oxygen.
*Some compounds like water (H2O) and ammonia (NH3) are always referred to by their common names.
REVIEW: Naming Binary Compounds
Three different types of binary compounds:
-Type I: Ionic compounds with metals that always form a cation with the same charge.
-Type II: Ionic compounds with metals (usually transition metals) that form cations with various charges.
-Type III: Compounds that contain nonmetals.
Flow chart for naming binary compounds:
Binary compound? If no, type III and they use prefixes. If yes, metal is present, does the metal form more than one cation? If no, type I, use the element name for the cation. If yes, type II, determine the cation charge; use a roman numeral after the element name.
Naming compounds that contain polyatomic ions:
-polyatomic ions are charged entities composed of several atoms bound together.
-polyatomic ions are assigned special names that you must memorize in order to name the compounds containing them.
Polyatomic Ions:
NH4+ is ammonium
NO2 - is nitrite
NO3 - is nitrate
SO3 2- is sulfite
SO4 2- is sulfate
OH- is hydroxide
CN- is cyanide
PO4 3- is phoshate
CO3 2- is carbonate
HCO3- is hydrogen carbonate (bicarbonate common name)
C2H3O2- is acetate
MnO4- is permanganate
Cr2O7 2- is dichromate
Oxyanions
Oxyanions are polyatomic anions that contain an oxygen and an atom of a different element.
-When there are two emembers in the series, the name of the one with the smaller number of oxygen atoms end in -ite, and the name of the one with the larger number ends in -ate.
-When more than two oxyanions make up a series, hypo-(less than) and per-(more than) are used as prefixes to name the members of hte series with the fewest and the most oxygen atoms, respectively.
Naming Ionic Compounds containing polyatomic ions:
Naming ionic compounds containing polyatomic ions is very similar to naming binary ionis compounds.

When a metal is present that forms more than one cation, a roman numberal is required to specify the cation charge, just as in naming type II binary compounds. EX: Fe(NO3)3 is named iron (III) nitrate.

When more than one polyatomic ion appears in a chemical formula, parentheses are used to enclose the ion and a subscript is written after the closing parenthesis.

*some examples can be found in the book on pages 138-140*
Naming Acids
-Acids are molecules, when dissolved in water, produce H+ ions (protons) and may be viewed as a molecule with one or more H+ ions attached to an anion.
-The rules are different for naming acids depending on whether the anion contains oxygen.
Rules for naming acids:
1. If the anion does not contain oxygen, the acid is named with the prefix hydro- and the suffix -ic attached to the root name for the element. EX: hydrochloric acid (HCl) and hydrocyanic acid (HCN)
2. When the anion contains oxygen, the acid name is formed from the root name of the central element of the anion or the anion name, with a suffix of -ic or -ous. When the anion name ends in -ate, the suffix -ic is used. EX: H2SO4 is sulfuric acid and HC2H3O2 is acetic acid (anions end in -ate) H2SO3 is sulfurous acid and HNO2 is nitrous acid.

*NEED TO KNOW THIS*
Flow for naming acids
Does the anion contain oxygen? No, hydro- plus anion root plus -ic (hydro-anion root-ic acid). Yes, check the ending of the anion name. If -ite, anion or element root plus -ous; (root)ous acid. If -ate, anion or element plus -ic; (root)ic acid.

Examples on page 40-41
Writing formulas with names:
The process can be reversed when the formula is written from the chemical name.
-You must learn the name, composition, and charge of each of the common polyatomic anions and the NH4+ cation.
-You must also learn the names of the common acids (hydrochloric acid, sulfuric acid, nitric acid, phosphoric acid, and acetic acid).
SIGNIFICANT FIGURES, con't.
3. Exact Numbers: Exact numbers are numbers determined from counting or from definitions and are assumed to have an unlimited number of significant figures. (Ex: 10 experiements, 8 people, 1 inch = 2.54 cm). They do no limit the number of significant figures in a caculation.
SIGNIFICANT FIGURES IN SCIENTIFIC NOTATION:
-The number 100. written in scientific notation is 1.00 x 10 to the 2nd (and has 3 SF).
-The number 0.000060 written in scientific notation is 6.0 x 10 to the -5th power (and has 2 SF's).
Elements to know: part two
Mg - Magnesium
Hg - Mercury
Ni - Nickel
N - Nitrogen
O - Oxygen
P - Phosphorus
Pt - Platinum
K - Potassium
Ag - Silver
Na - Sodium
S - Sulfur
Sn - Tin
Zn - Zinc
COMMON SIMPLE ANIONS, con't: MEMORIZE
ANION BINARY COMPOUNDS:
Hydride=H-
Fluoride=F-
Chloride=Cl-
Bromide=Br-
Iodide=I-
Oxide=O2-
Sulfide=S2-

*Group 6=sulfer group=2- ion*