Chem I Final Review

Ionic Compound

Compound where a metal reacts with a nonmetal

Polar Covalent Bond

Unequal sharing of electrons

If Electronegativity = Zero

Bond is covalent

If Electronegativity = Intermediate

Bond is polar covalent

If Electronegativity = Large

Bond is ionic

Dipole Moment

When molecules have a center of positive charge and a center of negative charge

aka when, in a compound, one element has a higher electronegativity than the other

Ion size determines…

Electronegativity increases…

Right way

Electronegativity decreases…

Down way

Cation formation occurs…

On left of periodic table

• electrons occupy high energy orbitals, so can lose electrons high energy electrons but not low —> limits size of positive charge on cations

Anion formation occurs…

On right of periodic table

• electrons occupy low energy orbitals

To create cations…

For main group elements, take electron from highest energy orbital

For transition elements, take uppermost s subshell first before valence d electrons

Cations Size

Smaller than atom formed from

Anions Size

Larger than atom formed from

Bond breaking is a process of

Endothermic process

Bond formation is a process of

Exothermic process

Calculating ΔH for a given reaction

= energy required to break bonds – energy released when bonds form

Calculating Formal Charge

Atom’s group # - [Amount of lone pair electrons + ½ amount of bonding pair electrons]

Electron-Pair Geometry: 2 groups around center

Linear

Bond angle: 180

Electron-Pair Geometry: 3 groups around center

Trigonal-planar

Bond angle: 120

Electron-Pair Geometry: 4 groups around center

Tetrahedral

Bond angle: 109.5

Electron-Pair Geometry: 5 groups around center

Trigonal-bipyramidal

Bond angle: 120 and 90

Electron-Pair Geometry: 6 groups around center

Octahedral

Bond angle: 90

If sum of lone pairs and atoms is 2, the hybridized orbital is…

Two sp orbitals

Unhybridized orbitals: Two p orbitals

If sum of lone pairs and atoms is 3, the hybridized orbital is…

Three sp² orbitals

Unhybridized orbitals: One p orbital

If sum of lone pairs and atoms is 4 the hybridized orbital is…

Four sp³ orbitals

Unhybridized orbitals: None

Geometry of sp hybrid orbital

Linear

Geometry of sp² hybrid orbital

Planar-trigonal

Geometry of sp³ hybrid orbital

Tetrahedral

When 1s orbitals combine…

They form two sigma molecular orbitals

• one bonding, one antibonding

When 2p orbitals combine…

They do

1. Direct head-to-head to form sigma and sigma⁺ molecular orbitals

2. Indirect, parellel to form pi and pi⁺ molecular orbitals

Molecular Geometry of 3 atoms + 0 lone pairs

Trigonal Planar

Molecular Geometry of 2 atoms + 1 lone pairs

Bent

Molecular Geometry of 4 atoms + 0 lone pairs

Tetrahedral

Molecular Geometry of 3 atoms + 1 lone pairs

Trigonal Pyramidal

Molecular Geometry of 2 atoms + 2 lone pairs

Bent

Molecular Geometry of 6 atoms + 0 lone pairs

Octahedral

Molecular Geometry of 5 atoms + 1 lone pairs

Square pyramidal

Molecular Geometry of 4 atoms + 2 lone pairs

Square planar

Molecular Geometry of 5 atoms + 0 lone pairs

Trigonal bipyramidal

Molecular Geometry of 4 atoms + 1 lone pairs

Seesaw

Molecular Geometry of 3 atoms + 2 lone pairs

T-shaped

Most likely Resonance is…

The one with the smallest formal charge

Paramagnetism

Presence unpaired electrons

Diamagnetism

Presence of paired electrons

Bond order

(# of bonding electrons - # of antibonding electrons) / 2

Larger bond order

Greater bond strength

Steps for Drawing MO Orbital Energy Diagram

Think of it as combining two atoms of the same element: Atom_A and Atom_B

1. Draw out left to right the amount of electrons in the 1s orbital for the element on both sides

2. Draw the bonding sigma 1s orbital and antibonding sigma 1s orbital

3. Transfer electrons from both sides into bonding sigma 1s orbital first, then antibonding if needed

4. Repeat for sigma 2s, pi 2p, etc

Hydrogen Bonding

When hydrogen reacts with N, O, F

Wavelength (λ)

The distance between two peaks or troughs in a wave.

Electromagnetic Radiation

Radiant energy that exhibits wavelike behavior and travels at the speed of light.

Frequency (v)

The number of waves or cycles per given time period that passes a given point in space.

Photon

A stream of individual packets of energy associated with electromagnetic radiation.

Quantum

A packet of energy; energy is transferred only in whole multiples of hv.

Photoelectric Effect

The emission of electrons from a metal when light strikes it.

• light must have a minimum frequency for electrons to be emitted

Threshold Frequency (V_o)

The minimum frequency required to remove an electron from a metal's surface.

• Greater than threshold frequency = light emits

• Lesser than threshold frequency = no light

Formula for Finding Kinetic Energy of an Electron

KE_electron = ½ mv² = hv -hv_o

• m - mass of electron

• v - velocity of electron

• hv - energy of incident photon

• hv_o - energy req to remove electron from metal’s surface

Bohr’s Theory of Hydrogen Atom

a) Has quantized energy levels

b) ELectrons only have certain allowed circular orbits

c) More tightly bound electron = more energy becomes negative

• energy released from system as electron brought closer to nucleus

Limitation: Electrons don’t move in circular orbits

Orbital

A specific wave function that describes a region of space where an electron is likely to be found.

Heisenberg Uncertainty Principle

The more accurately we know an electron's position, the less accurately we can know its velocity, and vice versa.

Excited State

A state of an atom in which electrons have absorbed energy and moved to higher energy levels.

Quantized Energy Levels

Discrete energy levels that electrons can occupy within an atom.

Energy Change (ΔE)

The difference in energy between the final state and initial state of an atom.

Radial Probability Distribution

The measure of the likelihood of finding an electron in a particular region of space around the nucleus.

Degenerate Orbitals

Orbitals that have the same energy level.

Noble Gas Configuration

A shorthand notation that uses the last noble gas to denote electron configuration.

Waves Have Three Characteristics:

Wavelength, frequency, and speed

Long wavelength

Short frequency

Short wavelength

Long frequency

What is the relationship between wavelength and frequency?

They are inversely related.

Formula Relating Wavelength, Frequency, and Speed of Light

c = λv

• c - speed of light (2.9979 × 10⁸ m/s)

• λ - wavelength (m)

• v - frequency (cycles per second)

Converting wavelength from nm to m

Multiply by 10^-9 m aka use conversion factor 10^-9m / 1 nm

Calculating For Energy of a Photon

E_photon = hv = hc / λ

• h - Planck’s constant (6.626 × 10^-34 J*s)]

• v - frequency of radiation

• λ - wavelength of radiation

Wavelength of Gamma Rays

10^-12 m

Wavelength of X-rays

10^-10 m

Wavelength of Ultraviolet

10^-8 m

Wavelength of Visible Light

4 × 10^-7 to 7 × 10^-7 m

Wavelength of Infrared Light

10^-4 m

Wavelength of Microwaves

10^-2 m

Wavelength of Radio Waves

1 to 10⁴ m

Relationship Between Energy and Wavelength

They are directly related.

Dual Nature of Light

Light can be a wave and photons

Calculating Wavelength for a Particle

de Broglie’s Equation

λ = h / mv

• h in units of J*S or kg*m²/s

1. Look @ units given and determine units of h to cancel out other units

2. Check units of mass and frequency and determine if conversions are needed

3. Plug into eqn and cancel out units

Calculating Energy Required to Move Electron from a Level to Another

Note: Removing electron from its ground state means n_initial = 1 to n_final = ∞

How to Calculate Wavelength of Emitted Photon

Quantum Model for Hydrogen Atom

A model where the electron is assumed to behave as a standing wave

• electron moves in certain allowed circular orbits

• assumes angular momentum of electron happens in certain increments

Quantum Numbers

Series of #s that describe properties of orbitals

Principal Quantum Numbers (n)

Tells energy level the electron occupies

• As n increases:

  1. Orbital becomes larger

  2. Electron gets farther away from nucleus

• Value of n correlates to # that goes before orbital sublevel letter (s, p, d, f)

Angular Momentum Quantum Number (ℓ)

Tells type of orbital we’re in

• = 0→ s

• = 1 → p

• = 2 → d

• = 3 → f

Has values from 0 to n-1

Magnetic Quantum Number (m_ℓ)

Tells the orientation of an orbital in relation to other orbitals in the same subshell

Has values from -ℓ to ℓ, including 0

Wavefunction Nodes (Nodal Surfaces)

Areas of zero probability in an orbital

• # of nodes increases as n increases

p Orbital Shape

Dumbbell / peanut shape

ℓ = 1

d Orbital Shape

Clover leaf shape

ℓ = 2

s Orbital Shape

Spherical shape

• as n increases, so does size

ℓ = 0

Number of Nodes for s Orbitals is Given by

n - 1

ℓ Angular Nodes

Nodes that cut through the nucleus

Radial Nodes

Nodes that circumscribe the nucleus

• given by n - 1 - ℓ

Valence Electrons

Electrons present in outermost principal quantum level (n) of an atom

To find:

1. Identify largest n value (i.e. 4p⁶)

2. Count the amount of electrons in that specific level; these are valence electrons

Core Electrons

Inner electrons, aka not valence electrons

Periodic Trends of Atomic Radius

Decreases right way

Increases down way