Chem 1 (College) Final Study Guide

Atomic Mass

The weighted average of all naturally occurring isotopes of the element

Protons

Atomic mass - neutrons

Neutrons

Atomic mass - protons

Electrons

Protons ± charge

*negative charge = more electrons

*positive charger = less electrons

Electron shells (n)

A group of atomic orbitals representing a specific energy level where electrons are found. While numbers greater than 0 (1-7).

Electron Sub-shells (cursive l)

Subdivisions of a shell (s, p, d, f)

S-orbitals

2 electrons

P-orbitals

6 electrons

d-orbitals

10 electrons

f-orbitals

14 electrons

Coulumb’s Law

F=k(q1•q2)/(r2)

K=constant

(q1•q2)=particle charge

r=distance between particles

F=attractive/repulsive force

Orbital Chart

Always fill entirely before moving on and fill a degenerate before pairing

Order to fill orbitals

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

Electron configuration

Shortened notation of the electronic structure of an atom or ion

Steps:

1. Determine total number of electrons involved

2. Start with lowest energy orbitals (1s) and add two electrons to each orbital till you run out

3. Double check that your used electrons = the total electrons

Electron configuration example (O)

Oxygen : 1s² 2s² 2p^4

Ion electron configuration

Anions = add extra electron to next available orbital

Cations = remove electron from the highest energy orbital

Noble Gas Abbreviations

[He] - 1s²

[Ne] - 1s² 2s² 2p^6

[Ar] - [Ne] 3s² 3p^6

[Kr] - [Ar] 4s² 3d^10 4p^6

[Xe] - [Kr] 5s² 4d^10 5p^6

[Rn] - [Xe] 6s² 4f^14 5d^10 6p^6

Effective Nuclear Charge (Zeff)

The pull on valence electrons by the nucleus.

Major: decreases from period 1-7

Minor: increases from group 1-18

Atomic Size Trend

The diameter of the element in its neutral state

Major: increased from period 1-7

Minor: decreases from group 1-18

Ionic Size Trend

The diameter of an element in its charged state

Cations: size decreases with each electron removed

Anions: size increases with each electron added

Isoelectronic

Same electron configuration but different element.

Ex. Na+, Ne, F- (all have 10 electrons)

Electronegativity (EN)

An atoms tendency to attract electrons to itself. High EN=want electrons, low EN=doesn’t want electrons

Major: decreases from period 1-7

Minor: increases from group 1-17

Formula mass

The mass of 1 formula unit of a compound.

Atomic mass of element x number of atoms in the formula

Ex: C6H12 = (6•12.01)+(12•1.008) = 84.156 amu (atomic mass unit)

Ionic bonds

Electrons are exchanged between atoms

Cations +

Anions -

Polar covalent bonds

Electrons are shared between atoms unevenly

Covalent bonds

Electrons are shared between atoms equally

Change in EN

EN(atom 1) - EN (atom 2)

change in En < 0.4 its pure covalent

1.8 > change in EN > 0.4 its polar covalent

Change in EN >1.8 (or it’s a metal) its ionic

Naming Cations

Polyatomic: Ammonium (NH4+)

Monatomic: name of element + Roman numerals (if charge is variable)

• I, II, III, IV, V, VI, VII, VIII

• Group 1 (+1), group 2 (+2), Ag (+1), Zn/Cd (+2), Al/Sc (+3)

Naming Anions

Polyatomic: later

Monatomic: base name + ide (Chlorine —> Chloride)

• fixed charges: Halogens (-1), group 16 (-2), group 15 (-3), group 14 (-4)

Octet Rule

Elements wants to hold onto 8 valence electron to be stable

Exceptions:

• H (only needs 2)

• Any atom below period 2 (can have more than 8)

Lewis Dot Structures

1. Determine total number of valence electrons

2. Arrange atoms (least EN atom in the center unless H or specified)

3. Add VE to each atom (dots)

4. Draw bonds between atoms (lines)

5. Check if every atom has a full octet and total VE matches ones used.

*if ion, add an electron for anions and take away one for cation

Formal Charge

A hypothetical charge an atom would have if the bonding electrons were equally shared

FC = (# of VE) - (# of non bonding electrons) - (1/2 # of bonding electrons)

Resonance Structure

Has the same chemical formula, same connectivity of atoms (formation), but a different arrangement of electrons in the compound

Isomer

Has the same chemical formula, but a different connectivity of atoms (formation) and different arrangement of electrons.

Valence Shell Electron Pair Repulsion Theory (VSEPR)

Electron groups will adopt a geometry with the largest possible angle between electron groups. Electron groups being one lone pair, one single bond, one double bond, one triple bond.

Electronic Geometry

Treat all electron group identically.

Linear, trigonal plantar, tetrahedral, trigonal bypyramid, and octahedral

Molecular Geometry

Differentiates between bonding and nonbonding electron groups

*has lone pairs

Sigma Bond

Always one sigma bond in every atom

Pi Bond

Only always between 2 p-orbitals

Hybridized Orbitals

sp orbitals, sp² orbitals, sp³ orbitals

• orbital = 1s + 1p = 2sp orbital

• sp² orbital = 1s + 2p = 3sp² orbital

• sp³ orbital = 1s + 3p = 4sp³ orbital

*orbitals put in = orbital put out

*look for number of electron groups 2 groups = sp, 3 groups = sp², 4 groups = sp³

Line angle drawings

A method of drawing structures where C and H are implicit. C goes on the end of every line and H goes on C to fill an octet unless otherwise specified.

Net dipole moment

Overall measure of a molecules polarity

Polar molecules must

1. Have at least one polar covalent bond

2. Have an overall net dipole moment (doesn’t cancel out)

Intermolecular forces (IMFs)

Attractive forces between molecules

Ex: dispersion forces, dipole-dipole forces, and hydrogen bonds

Dispersion forces (LDFs)

Arise from temporary dipoles in molecules. If the molecule has electrons, it has dispersion forces.

Dipole-Dipole forces (Dip-Dip/Dip²)

Arise from permanent dipoles in molecules aligning. Must be a polar molecule. In general, string than LDFs

Hydrogen Bonding (H-Bonds)

A subset of dip² forces. Molecule must contain a H bonded to N, O, or F. Externally strong bonds

Boiling point

Liquid to gas. A higher/stronger IMF means a higher boiling point and vise versa.

Melting point

Solid to liquid. A higher/stronger IMF means a higher melting point and vise versa

Viscosity

Resistance to flow (thickness). A higher/stronger IMF means a higher viscosity and vise versa.

Surface tension

Liquid state. A higher/stronger IMF means a higher surface tension and vise versa

Solubility

Dissolving. Molecules with the same/more similar strength IMF dissolve each other

Mass to moles

Molar mass

Moles to objects

Avogadro’s number (6.022×10²³)

Molarity (M)

M = moles/liters

Also concentration ( [L] )

Keq formula

[products]^# / [reactants]^#

only (aq) and (g) reactions

Keq numbers

Keq>1 products>reactants

Keq<1 products<reactants

Keq=1 products=reactants

Acids

HCl, H2, Se, H20, NH4+, Hbr, HNO3, etc.

Carboxylic Acid

Alcohol Acid

Thiol Acid

Bases

OH-, NH3, C2H3O2-, H20, etc.

Primary Amine

Secondary Amine

Tertiary Amine

Strong acids/bases…

completely dissociate in water

Weak acids/bases…

dont completely dissociate in water

pH formula

-log[H3O+]

[H3O+]

10^(-pH)

Percent Ionization

[H3O+] / [HA] x100

Ka formula

[A-] x [H3O+] / [HA]

weak acid = Ka<1

strong acid = Ka>1

pKa formula

-log(Ka)

Kb foruma

[OH-] x [HA] / [A-]

weak base = Kb<1

strong base = Kb>1

pKb formula

-log(Kb)

Kw formula

[OH-] x [H3O+]

1×10^14

pKw

14

pKa + pHb

Keq>Q

makes more products

Keq<Q

makes more reactants

Keq=Q

at equilibrium

Exothermic

produces heat and increases the temp of their surrounding, heat is a product

Endothermic

absorbs heat and decreases the temp of their surroundings, heat is a reactant