Exam 3 - Gen Chem 107

energy

the capacity to supply heat or do work

potential energy

• bonds

• stored energy

speed of hot particles

fast

speed of cold particles

slow moving

thermal equilibrium

when one hot thing touches a cold thing, the kinetic energy transfers and thermal equilibrium will occur and they will start moving at the same pace

potential energy equation

mgh (mass x acceleration due to gravity)

kinetic energy equation

½ mv² (mass and velocity)

law of conservation of energy

energy cannot be created or destroyed, only transferred or transformed

aspects that affect heat transfer

• size

• starting temp of each substance

• identity of materials

formula for heat energy

q = mcΔT

• q = heat energy in joules

• m = mass

• c = identity

• ΔT = change in temperature

heat capacity (C)

quantity of heat it absorbs or releases when it experiences a temperature chance of 1 degree celsius

• extensive property

specific heat capacity (c)

the quantity of heat required to raise the temperature of 1 gram of a substance by 1 degree celsius

• intensive

if a substance gains thermal energy (Tfinal-Tinitial), the value of q =

positive

if a substance loses thermal energy (Tinitial-Tfinal), the value of q =

negative

systems + surroundings =

universe

surroundings

the water the reaction takes place in

• this is what we can monitor

system

the reaction itself

exothermic process

a change where system releases heat

endothermic process

a change where a system absorbs heat

skeletal equation of endothermic reactions

reactants + energy → products

skeletal equation of exothermic reactions

reactants → products + energy

Closed system

Surroundings get hotter, system’s temp decreased OR surroundings get cooler and system’s temp increased

In calorimetry… heat lost + heat gained =

0

Internal energy, U

= q + w

Enthalpy, H

• heat only under constant pressure

• In KJ/mole

= qp (subscript of p)

ΔH < 0

Exothermic

ΔH > 0

Endothermic

State function

Property or condition that relies only on the current condition, or state, that the object or system is in. How the object or system achieved that condition is reach is not relevant

Examples of state functions

• temperature

• Enthalpy

• Pressure

• Volume

Standard enthalpy of formation of an element in its most stable form

0 under standard conditions

Calculating ΔH for a reaction

ΔH = sum of ΔH(products) - sum of ΔH(reactants)

Hess’s law

If a process can be written as the sum of several stepwise processes, the enthalpy change of the total process equals the sum of the enthalpy changes of the various steps

Exothermic process turned around

Becomes endothermic

Conventions of thermochemical equations

• If the coefficients of the chemical equation are multiplied by some factors, the enthalpy must also be multiplied by the same factor

2 perpendicular oscillating fields

Electric and magnetic

Wavelength

Distance from peak of one wave to peak of another wave

• units of length

Frequency

How many wave forms are going to pass through in one second

• units of hertz

Amplitude

How tall the peak is or how low the trough is

• Independent of wavelength and frequency

Energy and frequency

Increase as you go up the light list

Wavelength (in accordance to the light lost)

Decreases as you go up the list

Harmonics

Integer multiples of the original ½ wavelength

Nodes

Divot the middle of the waves

One wave form

2 peaks and 2 troughs

Relationship between wavelength and frequency

Wavelength x frequency = speed of light ( c )

Radial nodes

Circles

Angular nodes

Radii

How we know light behaves like a wave

Diffraction as light passes through slits

Diffraction pattern

Some waves’ peaks combine and form an extra strong power, when the peak and trough interact, they cancel each other out

• basis of classical mechanics

Experiments /observations that do not fit “light is a wave”

• blackbody radiation

• Photoelectric effect

• Line spectra of atoms

Blackbody radiation

• light from sun gives us a continuous spectrum

• Approximated by blackbody radiation

• Light given off by an object heated until it glows

• The intensity should have kept rising, but it fell off, and that didn’t fit

Max Planck equation

E=hv = hc/wavelength

Particle nature of light

Photons

Albert Einstein’s first Nobel prize

Photoelectric effect

Photoelectric effect

Phototube with two electrodes that are not touching - closed circuit

• when they shine light into the phototube, it produces a measurable current

• Couldn’t explain that when using the purple/violet light they could see the current, but when they used red light they could not

• Purple light had enough energy to knock the electron loose and produce the current

• Einstein used collision to explain this and described light as a packet - energy depends on frequency or inversely on wavelength

Line spectra of atoms

If we have a source of white light that travels through and is split with a prism, we will get all the types of light within the white light

Bohr Model

• occupy discrete energy levels and only exist in the available energy levels

• May move between energy levels by either absorbing or emitting energy

• Have energy levels designated by specific value for n

What the Bohr model gets right

• electrons move between discrete energy levels

• Accompanied by the absorption or emission of light whose energy corresponds to the difference in energy levels

Limitations of Bohr Model

• works for H, and He, but not any other element

• Relied on a precise orbit for the electron that is not correct

Electrons are…

Waves!

Wave function

3 dimensional stationary wave, psi

Psi²

Probability of finding an electron = orbital

• can be used to determine the distribution of the electron’s density with respect to the nucleus in an atom

De Broglie Wavelength

Wavelength of matter (particles)

Wavelength of matter

h/mv = h/p

• H = plank’s constant

• M = particle mass

• V = particle velocity

• P = particle momentum

Heisenberg uncertainty principle

Use of wave functions

Can be used to determine the distribution of the electron’s density with respect to the nucleus in an atom

Quantum Theory of Electrons in Atoms

• electrons in atoms can exist only in discrete energy levels but not between them

• Energy of an electron in an atom is quantized

• Energy can be equal only to certain specific values and can jump from one energy level to another, but not transition smoothly or stay between the levels

Quantum numbers

Mathematical solutions to Schrödinger equation

• no 2 electrons in an atom have all the same quantum numbers

Principle quantum number, n

• aka shell

• Defines energy level

• Size of the orbital

• Describes how far away from the nucleus an electron is

Angular momentum, l

• aka shape of the orbital

• l = 0, 1, 2,

• Max is n-1

l=0

S orbital

l=1

P orbital

l=2

d orbital

l=3

f orbital

S subshell

Spherical shape

P Subshell

Dumbbell shape

D and f orbitals

More complex

Orbital shapes

Represent the 3 dimensional regions where the electron is likely to be found

Radial nodes formula

n - l - 1

Nodes in a 1s orbital

1-0-1=0

Nodes in a 2s orbital

2-0-1=1

Nodes in a 3s orbital

3-0-1=2

Magnetic quantum number, m1

Orientation of the orbital in space

• -l, … 0 …, +l

Spin quantum number, ms

Intrinsic property of electron

• not a solution of wavefunction

2 possibilities for spin number

+1/2 (up) or -1/2 (down)

Pauli exclusion principle

• an electron in an atom is completely described by 4 quantum numbers

• No two electrons in the same atom can have exactly the same set of all four quantum numbers

Degenerate

Have the same energy

• All energy units within 3p and 2p are the same energy

Shielding

Repulsion from inner electrons

• electrons are attracted to nucleus but they are repulsed by each other

Effective nuclear charge

Z - shielding

Outer electrons have higher/lower energy

Higher - they are further from the nucleus

In any atom with 2 or more electrons…

The energy of the orbitals increase within a shell in the order s<p<d<f

Aufbau principle

Fill lowest energy orbitals first

Pauli exclusion principle

No two electrons have the same QN - 2 electrons per orbital

Electron configuration

The arrangement of electrons in the orbitals of an atom

3 pieces of information in an electron configuration

1. The principal quantum shell, n

2. The letter that designates the orbital type (l)

3. A subscript number that designates the number of electrons in that particular shell

Orbital diagrams

Pictorial representations of the electron configuration, showing the individual orbitals and the pairing arrangement of electrons

Hund’s rule

The lowest energy configuration for an atom with electrons within a set of degenerate orbitals is that having the maximum number of unpaired electrons

Two periodic table columns that have abnormal electron configurations

Group 6 and group 11

Electron configurations of anions

Electrons are added in the order predicted by the aufbau principle

Electron configuration of a cation

Add in electrons for neutral atom, then remove electrons

• transition metals: remove a electrons before d