Chem Quiz 1

Question 1

Dalton’s Atomic Model (1809)

Answer 1

Matter consists of tiny atoms
Atoms are indestructable
All atoms of one element are identical
Atoms of diff elements differ in mass and properties
Atoms combine in whole number ratios

Question 2

Daltons Atomic Model Limitations

Answer 2

Couldnt explain why atoms of diff elements combine in the ratio in which they do
Science in 1850s and 1900s suggested atoms not divisible

Question 3

JJ. Thompson - Cathode Ray Experiment

Answer 3

Discovered negatively and positively charged particles using cathode ray experiment
Ray was produced from negative end towards positive end, and was deflected by electric feild
Thompson hypothesized that ray in cathode ray was composed of stream of negatively charged particles

Plum pudding model: electrons are negative particles embedded in a positive ball
Limitation: later disproved by Rutherford’s gold foil experiment

Question 4

Robert Milkman (1909) - Electron

Answer 4

Measured charge of an electron using oil drop experiment (in coulombs)
Also electron mass

Question 5

Rutherford (1909) - Gold Foil Experiment

Answer 5

Rutherford shot alpha particles (positively charged) at thin sheet of gold foil
1/8000 particles were deflected significantly, some 90 degrees or more
Rutherford suggested that these deflections were due to particles encountering an intense electric feild at centre of the atom
Atoms are mainly empty space
Tiny positively charged nucleus that contains most of atoms mass

Question 6

Limitations of Gold Foil Experiment (Rutherford)

Answer 6

(1) Positive charges in nucleus should repel each other and break apart nucleus. Doesn’t explain total mass of atom –> Solved when Chadwick discovered neutron in 1932!

(2) An electron in motion should give off a continous spectrum of light energy, in doing so, will lose energy, orbital radius will decrease an eventually spiral into the nucleus, destroying the atom –> doesn’t happen!

Question 7

Atomic Spectra

Answer 7

Visible Light is part of the elecromagnetic spectrum. EM radiation travels at the speed of light

As wavelenght increases, frequency decreases

Classicial physics would predict that energy (light) emitted from excited electrons should be continous. Its not! When electrons absorb energy, a pattern of discrete lines is observed

Question 8

Black Body Radiation

Answer 8

Black body is a body that emits or absorbs EM radiation of all wavelenghts (ex. sun)

As you heat a blackbody, it will emit radiation of decreasing wavelenght (higher frequency and energy) dependent on temperature. At 4000K and above, hot objects begin to glow.

Classical physics predicts that the intensity would continue to rise the more the object is heated and would approach to infinity as the wavelenght decreases

Question 9

Photoelectric Effect

Answer 9

When light (high energy) shines on a metal, electrons are emitted from the surface of the metal

As f increases, the kinetic energy of the electron increases

As amplitude increases, current increases

Question 10

Einstein (1921 Nobel Prize) - Photons

Answer 10

Light sometimes behaves as particles, photons (quantized units). Light photons interact with metal atoms on the surface and if energy is sufficieint, will cause electron emission.

Question 11

Bohr Model (1913)

Answer 11

Atoms have only specific energy levels (with specific radii) that electron can occupy

Within an energy level, electrons dont emit energy, they are at a steady energy state

Electrons change levels by absorbing or emitting energy (light). The amount of energy exists in disecrete quantities and is not continous

Max Planck called these packets of energy, Quantums
In 1905, Einstein determined that light is also quantized
Photons: particle-like packets of quantized energy

All EM energy travels in the form of photons of energy

Question 12

Emission Spectrum

Answer 12

Spectrum of electromagnetic radiation emitted by an atom; results when atom is returned to lower energy state from higher energy state.

blackground=black, think colored lines.

Question 13

Absorption Spectrum

Answer 13

Spectrum of electromagnetic radiation absorbed by an atom; results when atom jumps from lower to higher energy state

background=colored, thin black lines

Question 14

Limitations of Bohrs Model

Answer 14

Bohrs Model explained one electron atoms, not atoms with more than one electron… ex. H, He+, Li2+, Be3+
Couldnt explain emission spectra of atoms with many electrons
Mercury for example, has smaller spaced between colored lines - This suggested that there were smaller energy differences between energy levels
A new model was required to explain the spectra of all types of atoms

Question 15

Louis de Brogile (1924) - Wave Particle Duality

Answer 15

Wave Particle Duality - Matter have wave like properties

Combined Planks equation and Einsteins equation - replaced speed of light with speed of particle

Wavelenght of basketball is unimaginably small, it has no observable effect in the baseballs motion

Experimental evidence = streams of moving electron produced diffraction patterns simillar to those produced by light waves

Question 16

Heisenberg (1927) Uncertainty Principle

Answer 16

On the tiny scale of atom, particle model of electron doesnt describe its properties - tends to act more like water wave than billard ball

Uncertainty prinicple: due to wave nature of matter, its impossible to predict both the position and monumentum of an electron with certainity

If one range, lets say position is known more precisely, the range in monumentum is larger and thus cannot be known precisely

for an electron, mass is very small so you cannot know the position or speed with much precision

Question 17

Schrodinger (1926)

Answer 17

Combined de Brogils ideas of matter waves with Einsteins idea of quantized particles (photons) and derived an equation that tells us probability that an electron is at a particular point

Solution to this equation gave rise to set of mathematical expressions called wave function –> has 3 variables (n, l, ml).. known as quantum numbers

When you substitude specific combinations of integers for each of these variables into the wave function, many solutions result. Each solution describes the 3D region in space around nucleus of an atom. Chemists call this region of space thats related to the wave function an atomic orbital.

Summary IMPORTANT:
1. Wave functions (solutions to wave equation) describe the electron in an atom

2. Square of wave functions: the mathematical probability of finding an electron in a certain region of space
3. In quantum mechanics, the electron probability density is all thats known
4. Orbital is the region around nucleus where theres a high probability of finding an electron

Question 18

Electron Orbitals

Answer 18

Electrons resemble standing waves around nucleus

Certain orbitals exist if they correspond to whole wavelenghts (1, 2, 3, etc) of electron vibrations, not part of waves (ex. 1.5)

Question 19

First Quantum Number (n)

Answer 19

Principal Quantum Number (n)
Positive int (ex:1, 2, 3, 4, ..) and so on

Specifies the energy level and relative size of an atomic orbital - the higher the n value, the higher the energy level and size of orbital

Larger the orbital, higher the probability of finding an electron further from nucleus

Largest possible number of electrons in an energy level is 2(n^2)

Question 20

Second Quantum Number (l)

Answer 20

Describes orbital shape (called orbital shape quantum number) - “l”

Sometimes referred to as angular monumentum quantum number

Positive integer values: 0 to (n-1), n-1 being max value

If n=1, l=0 (one possibe sublevel)
If n=2, l= 0 or 1 (two possible sublevels)
If n=3, l=0, 1, 2 (three possible sublevels)

l=0, sublevel s
l=1, sublevel p
l=2, sublevel d
l=3, sublevel f

To identify an energy sublevel (type of orbital) combine “n” value and letter of orbital shape

Ex. n=3, l=0 called 3s sublevel
n=2, l=1 called 2p sublevel

Question 21

Third Quantum Number (ml)

Answer 21

Magnetic Quantum Number (ml)
Can be from -l to +l
Describes orientation of orbital in space

If l=0, ml=0
If l=1, ml= -1, 0, or 1
In general, for a given value of l, there are 2l+1 values for ml

Total number of orbitals for any energy level if n^2

Question 22

Fourth Quantum Number (ms)

Answer 22

Spin Quantum Number (ms)

Results from the particle like nature of electrons - Electrons can be thought of as spinning about their axis

This is proven experimentally, as electrons are shown to have magnetic dipoles as they’re spinning

Because electrons are not classical particles, they dont truly spin as we would imagine a spinning top

ms can be either +1/2 or -1/2

Question 23

Pauli Exclusion Principle

Answer 23

1. Only two electrons of opposite spin can occupy an orbital
2. No two electrons have the same 4 quantum numbers

Question 24

Aufbau Principle

Answer 24

1. The number of electrons in an atom is equal to the atomic number
2. When adding electrons to orbitals, start by filling the orbital with lowest energy, and work your way up

• Remember that each oribital can have a max of two electrons

Question 25

Hund’s Rule

Answer 25

1. Every orbital in a subshell is occupied by 1 electron before any pairing can occur
2. All electrons in singly occupied orbitals have the same spin

Question 26

Exceptions to electron configuration rule?

Answer 26

Cr and Cu
d^9 should be d^10
d^4 should be d^5
———-
f^13 should be f^14
f^6 should be f^12

d and f orbitals are more stable when full or half filled, take one from s^2 to fulfill!

Question 27

Atomic Radius

Answer 27

Measured by taking half the distance between identical nuclei in a crystal of a metal element or between identical nuclei in a single covalent bond

Question 28

Atomic Radius Trend?

Answer 28

atomic radius is determined by n
as n increases, electrons are farther from nucleus

size increases as you go down a group. there is an increased sheilding effect from inner electrons

Effective nucleur charge (Zeff) is the net force of attraction between electrons and the nucleus. As Zeff increases, electrons are more strongly attracted to nucleus
Zeff = atomic # - sheilding effect

Size decreases as you go across a period to the right (more protons, more Zeff). Same outer energy level (same n) - little change in sheilding effect.

Question 29

Ionization Energy

Answer 29

The energy needed to remove an electron from a ground state atom (to overcome attraction between electron and proton)

Smaller IE = easier to remove
Larger IE = harder to remove

Atoms with multiple electrons have multiple ionization energies

1st IE < 2nd IE < 3rd IE (as you remove electrons, atom becomes positive, harder to remove electrons from positive atoms)

Question 30

Trends in First Ionization Energy?

Answer 30

1st IE levels increase across a period due increased Zeff

As Zeff increases (and size decreases), electrons are harder to remove so IE increases

IE decreases going down because atomic radius increases, so its easier to remove electron

Question 31

Why the Blipps?

Answer 31

Groups 2 and 3, electron is removed from p orbital rather than s orbital - Electron is farther from nucleus. There is small amount of repulsion from s electron

Groups 5 and 6, electron is removed from a doubly occupied p orbital
There is electron repulsion that aids in the removal of electron

Question 32

Electron Affinity

Answer 32

The change in energy that occurs hen an electron is added to a gaseous atom (AKA the ability of an atom to accept an electron)

High EA = high negative # = easier it is to add an electron (negative value means energy is given off when electron is added)

Low EA = low negative # or positive # (positive value means energy is required to add an electron)

EA decreases as you add more electrons (due to net negative charge) 1st EA > 2nd EA > 3rd EA

Question 33

Electron Affinity Periodic Trends

Answer 33

EA decreases as you go down a group

EA increases to the right