Chemistry 1 Exam 2 (Chapter 9-10)

Front 1

Ionic Bond

Back 1

metals and non metals

electrons transferred (from metal to nonmetal)

their oppositely charged ions attract each other, lowering their overall potential energy

Front 2

Covalent Bonds

Back 2

nonmetals and nonmetals

electrons shared between both nonmetals

the shared electrons react with the nuclei of both atoms, lowering their potential energy

Front 3

Metallic Bonds

Back 3

metals and metals

electrons pooled, 'lattice pool'

the electrons aren't localized over a single atom, but delocalized over the entire metal; the positively charged metal atoms are then attracted to the sea of electrons holding it together

Front 4

Why do chemical bonds form?

Back 4

because they lower the potential energy between the charged particles that compose atoms

coulombs law states that when two atoms approach each other the electrons of one atom become attracted to the nucleus of the other

Front 5

Metals

Back 5

tend to have low ionization energies, meaning that their electrons are easy to remove

Front 6

Nonmetals

Back 6

tend to have negative electron affinities, meaning that they readily gain electrons

Front 7

Cation

Back 7

metal (loss of electrons)

Front 8

Anion

Back 8

nonmetals (gain of electrons)

Front 9

Lewis Structure

Back 9

represents valence electrons of an element with dots

Front 10

Octet

Back 10

an element having 8 valence electrons, completing their outermost shell

noble gasses

many elements bond in order to complete their octets, becoming more stable

Front 11

Octet Rule Exceptions

Back 11

Too few:
Hydrogen (1 electron, 1 bond = 2 valence)
Beryllium (2 electrons, 2 bonds = 4 valence)
Boron (3 electrons, 3 bonds = 6 valence)

Expanded Octets:
Sulfur (5 electrons, 5 bonds = 10 valence)
Phosphorus (6 electrons, 6 bonds = 12 valence)

Front 12

Chemical Bond

Back 12

the sharing or transfer of electrons to attain stable electron configurations for the bonding atoms

Front 13

Lattice Energy

Back 13

the energy associated with forming a crystalline lattice of alternating cations and anions from gaseous ions

the easiest way to calculate lattice energy is with the born-haber cycle

Front 14

Sublimation

Back 14

going from a solid to a gas

Front 15

Dissociation

Back 15

separating of a compound

Front 16

Born-Haber Cycle

Back 16

a hypothetical series of steps that represents the formation of an ionic compound from its constituent elements

the steps are chosen so that each step is known except for the last one, which is lattice energy

the change in enthalpy is also known (ΔH final)

Front 17

Born-Haber Cycle formula

Back 17

ΔH final= ΔH1+ΔH2+ΔH3+ΔH4+ΔH5

change in enthalpy= (sublimation + bond energy + ionization energy + electron affinity + lattice energy/formation energy)

*for bond energy don't forget to divide if needed*

Front 18

Trend in Atomic Radius

Back 18

increases as you go down and to the left of the periodic table

Front 19

Coulomb's Lattice Energy Law

Back 19

E = K(Qpos + Qneg)/r

E: potential energy
Qpos: charge of cation
Qneg: charge of anion
R: atomic radius/distance between ions
K: proportionality constant

(the more negative a lattice energy, the stronger it is)

Front 20

Lattice Energy Trends with Radius

Back 20

as you go down a column, the radius increases, meaning that the lattice energy will be less because the ions aren't as close and do not release as much energy when the lattice forms

Front 21

Lattice Energy Trends with Charge

Back 21

as the charges of the ions increase, the lattice energy increases

charge affects lattice energy more than radius

multiply the charges of each element involved, and the one with the largest negative total will have the strongest lattice energy

Front 22

Lattice Energy and Melting Points

Back 22

higher lattice energy, the higher the melting point

Front 23

Exothermic and Endothermic Properties of Lattice Energy

Back 23

lattice energy becomes more exothermic (more negative) with increasing charge, and more exothermic with closer radii

lattice energy becomes more endothermic (less negative) with decreasing charge, and more endothermic with increasing radius

Front 24

Exothermic

Back 24

gives off heat, transfers thermal energy from the system to the surroundings

loss of energy

Front 25

Endothermic

Back 25

needs heat supplied into system from the surroundings to begin

increase in energy

Front 26

Double Bond

Back 26

two atoms share two pairs of electrons

Front 27

Triple Bond

Back 27

two atoms share three pairs of electrons

Front 28

Covalent bonds are highly...

Back 28

directional, the attraction between the two bonded atoms is due to the sharing of electron pairs in the space between them

the interactions are between the molecules, aka intermolecular forces

Front 29

Ionic bonds are...

Back 29

nondirectional and hold together an entire array of ions

Front 30

Molecular Compounds vs Ionic Compounds

Back 30

molecular compounds have lower melting and boiling points than ionic compounds

Front 31

Problem with Lewis Structures Regarding Electron Sharing

Back 31

it shows that in covalent bonds the electrons appear to be equally shared, but that's not the case. Each element has a different electronegativity and therefor a stronger or weaker pull for those electrons

Front 32

Trend in Electronegativity

Back 32

increases as you go up and to the right on the periodic table, with F being the strongest electronegative

Front 33

Electronegativities

Back 33

the element further up and to the right is more electronegative, symbolized by 𝛿- meaning the pull of electrons will tend towards this element (higher electron density). Bears a partial negative charge.

the element not as far to the right and up is more electropositive, symbolized by 𝛿+, the pull of electrons is not in their favor (lower electron density). Bears a partial positive charge.

Front 34

Polar covalent bonds

Back 34

when elements of two different electronegativites bond together, giving an unequal sharing of electrons

has a positive and negative pole

Front 35

Determining Bond Types by their Electronegativity

Back 35

(0 - 0.4) → small electroneg diff → covalent bond

(0.4 - 2.0) → intermediate electroneg diff → polar covalent bond

(2.0+) → large electroneg diff → ionic bond

To determine the difference, subtract the more electronegative atom to the less electronegative atom and compare the value.

Front 36

Electronegativity Related to Atomic Size

Back 36

the larger the atom, the less ability it has to attract electrons to itself in a chemical bond

Front 37

How to calculate percent ionic character

Back 37

___pm (4.80D/100pm) = x (measured D/x)(100)= % ionic

Hint 37

Xx

Front 38

Percent Ionic Character

Back 38

the ratio of a bond's actual dipole moment to the dipole moment it would have if electrons were 100% transferred from one atom to another

in general, bonds with 50%+ ionic character are ionic bonds

Front 39

How to write a Lewis Structure

Back 39

1. more electronegative elements go on the ends, the least electronegative goes in the center

2. calculate total number of electrons in compound
(if has neg charge add 1 electron and if pos charge subtract and electron)

3. Fill up the atoms with electrons, aka dots, and if you have too many electrons, create more bonds until octets are full, some will have more or less than a complete octet

Front 40

Resonance Structures

Back 40

represents all the possible representations of a compound with double arrows in-between them

some are more favorable structures, the one with the least formal charges will be the most favorable

Front 41

Resonance Hybrid

Back 41

the actual structure of a compound with multiple structures is more intermediate with the positions of electrons in each different structure

Hybrids are represented with a dashed line over the bonds where different electrons are shared/bonded. One end may have a double bond and another doesn't and they intermediately switch between sides the bond is located.

Front 42

Resonance Stabilization

Back 42

the delocalizing of electrons, by having intermediate positions for these electrons, creating them more stable

Front 43

Formal charge

Back 43

a fictitious charge assigned to atoms in a lews structure

represents the charge it would have if all bonding electrons were shared equally between the bonded atoms

Front 44

How to calculate formal charge

Back 44

# of electrons it's supposed to have (group number) - actual amount of electrons directly touching it

Front 45

Rules for formal charges

Back 45

1. the sum of all neutral charges on a molecule must be zero
2. the sum of all the formal charges of an ion must equal that charge
3. small (or zero) charges on individual atoms are better than larger ones
4. when formal charges cannot be avoided, the negative formal charge should reside on the most electronegative atom

Front 46

Odd-Electron Species

Back 46

have an odd number of electrons when bonded
the N on NO2 can have 7

Front 47

Incomplete Octets

Back 47

don't fill up to exactly 8 electrons when bonded

Hydrogen (1 electron, 1 bond = 2 valence)
Beryllium (2 electrons, 2 bonds = 4 valence)
Boron (3 electrons, 3 bonds = 6 valence)

Front 48

Expanded Octets

Back 48

have too many initial electrons and therefore can make more than 4 bonds

Sulfur (5 electrons, 5 bonds = 10 valence)
Phosphorus (6 electrons, 6 bonds = 12 valence)

Front 49

Bond Energies

Back 49

the energy required to break one mole of the bond in the gas phase

they are always positive because it takes energy to break a bond

compounds with stronger bonds tend to be more stable, and less chemically reactive

Front 50

Bond Lengths and Strengths

Back 50

triple / double / single
shorter → longer
stronger → weaker

Front 51

Bond Energies to Estimate Enthalpy Changes

Back 51

average bond energies are used to estimate the enthalpy change of a reaction

Formula:
ΔHrxn= (sum of bonds broken) -( sum of bonds formed)

broken bonds = positive
formed bonds= negative

Front 52

The 'thermic' of a reaction

Back 52

exothermic- when weak bonds break and strong bonds forms

endothermic- when strong bonds break and weak bonds form

Front 53

Valence Shell Electron Pair Repulsion (VSEPR) Theory

Back 53

based on the ides that electron groups, ('lone pairs', single-triple bonds and singe electrons) repel one another through coulombic forces

the repulsions between electron groups on interior atoms of a molecule determine the geometry of the molecule

Front 54

VESEPR Trends

Back 54

increased VSEPR repulsion from lone pairs reduces bond angles

lp vs lp repulsion > lp vs bp repulsion > bp vs bp repulsion

Front 55

Predicting Molecular Geometry

Back 55

1. draw lewis structure
2. count number of bonds and lone pairs on central atom
3. use VSEPR to predict the geometry

Front 56

Electron (VSEPR) Geometry

Back 56

count all electron bonding pairs (double bonds count as one)

the lone pairs do not influence

Front 57

Molecular Geometry

Back 57

the non bonding, lone pairs will influence the shape now

Front 58

How to Represent Molecular Geometries on Paper

Back 58

hatched wedge: bond going into the page
solid wedge: bond coming out of page
straight line: bond in plane of paper

Front 59

Determining Polar Molecules

Back 59

At least 1 polar covalent bond + asymmetrical shape

if it is a non polar molecule as a whole, it will not be polar

Front 60

Colors of Polar Molecules in 3D Representations

Back 60

𝛿- redish-yellowish color
𝛿+ blueish-greenish color

use molecular geometry to know shape

Front 61

Dipole Moment

Back 61

dipoles point to the more electronegative atom

no net dipole moment: sum of dipoles = zero (do not point in universal direction)

dipole moment: point in a universal direction

Front 62

Hybridization

Back 62

mixing of two or more atomic orbitals to form a new set of hybrid orbitals

Front 63

Sigma Bonds

Back 63

directly connect

electron density is between two atoms

single bonds

tip to tip laying down, with two ends touching formation

Front 64

Pi Bonds

Back 64

electron density above and below plane of nuclei

single and double bonds

side by side standing up formation with each tip overlapping

Front 65

Single, Double and Triple Bonds

Back 65

Single: 1 sigma
Double: 1 sigma, 1 pi
Triple: 1 sigma, 2 pi

Front 66

Molecular Orbital Theory

Back 66

bonds are formed from interaction of atomic orbitals to form molecular orbitals

Front 67

Antibonding Molecular Orbitals

Back 67

higher energy and lower stability

symbolized by sigma*

Front 68

Bonding Molecular Orbitals

Back 68

lower energy and greater stability

symbolized by sigma

Front 69

Bond Order Formula

Back 69

(1/2)(electrons in bonding - electrons in antibonding)

lower bond order = less stable

Front 70

How to fill in molecular orbital charts

Back 70

fill in electrons on left and right following all the filling rules (2s first, then 2p, etc)

then fill in center molecule with bonding (at bottom) electrons going in before anti bonding (at top) (ADD THE TWO ELECTRONS ON BOTH SIDES TO FILL IN CENTER)