Unit 1 - Structure and Properties

Question 1

Principal Quantum Number

Answer 1

-indicates relative size and energy of atomic orbitals
-represents energy level, or shells, the electrons can occupy in an atom
-as n increases, orbital size increases
-max number of electrons per energy level: 2n^2

Question 1

Orbitals (definition, different quantum numbers)

Answer 1

represent probability clouds that show a statistical distribution of where the electron is likely found

Described by 4 quantum numbers:
-principal quantum #
-angular momentum (secondary) quantum #
-magnetic quantum #
-spin quantum #

Question 2

Angular Momentum (secondary) Quantum Number (definition, how many electrons held by each shape)

Answer 2

-energy sublevels are contained within the principal energy levels
-gives shape of the subshell (s,p,d,f)
-L = 0 up to (n-1)
s orbitals: L=0, 1 s orbital per subshell, can hold total of 2e-
p orbitals: L=1, 3 p orbitals per subshell, can hold total of 6e-
d orbitals: L=2, 5 d orbitals per subshell, can hold total of 10e-
f orbitals: L=3, 7 f orbitals per subshell, can hold total of 14e-
energy increases from s->f orbitals as they get further from nucleus

Question 3

Magnetic Quantum Number

Answer 3

-represents the orientation of the subshells
-has values of -L to +L

Question 4

Spin Quantum Number

Answer 4

-represents the spin of the electron
-has values of +1/2 or -1/2

Question 5

Heisenberg Uncertainty Principle

Answer 5

it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time

Question 6

Pauli Exclusion Principle

Answer 6

-No two electrons in an atom can have the same FOUR quantum numbers
-means only e- of opposite spin can occupy an orbital

Question 7

Aufbau Principle (both parts)

Answer 7

1. Place electrons into the orbitals in order of increasing energy level
2. Each subshell of orbitals in the same energy level must be completely filled before proceeding to the next orbital or series of orbitals

Question 8

Hund’s Rule (both parts)

Answer 8

1. Whenever electrons are added to orbitals of the same energy sublevel, each orbital recieves one electron before pairing occurs
2. When electrons are added singly to seperate orbitals of same energy level, the electrons must all have the same spin

Question 9

Ground State vs Excited State of electron

Answer 9

Excited state is when valence electron absorbs energy, commonly in the form of jumping from its ground state (original energy level or orbital) to an empty orbital of a higher energy shell

Question 10

Isoelectronic Ions

Answer 10

-Atoms and ions that have the same number of electrons
-elements coming before and after noble gases form isoelectronic ions by gaining or losing electrons until they have the same # of electrons as the noble gas

Question 11

Valence Shell Electron Pair Repulsion Theory/VSEPR (definition, key ideas)

Answer 11

Allows us to predict the geometry of molecules based on the repulsion of electron pairs. There are five basic geometries, and 10 subtypes for if lone pairs are present, total 15 geometries.

Key Ideas:
1. e- pairs (both bonding and lone pairs around central atom) repel each other electrostatically
2. The molecular shape is determined by the positions of the electron pairs when they are a maxium distance apart

Question 12

Types of Electron Pairs

Answer 12

-Bonding pairs or lone pairs (non-bonding pairs)
-If all e- pairs are identical, then the molecular shape is the same as the e- pair arrangement (CCl4)
-If all e- pairs are NOT identical, then e- pair arrangements is only an approximation of the shape of the molecule (CCl3F)
—Three reasons for this: lone pairs occupy more space, space occupied by bonding pairs related to EN of the atom, double and triple bonds occupy more space

Question 13

Ionic Bond

Answer 13

-Metal and non-metal atom
-TRANSFER of electrons from metal to non-metal
-cation and anion
-ΔEN is 1.7 and above

Question 14

Polar Covalent Bond

Answer 14

-electrons are shared UNequally because 1 atom attracts them more strongly than the other atom
-dipole created, region of slightly positive charge, region of slightly negative charge
-ΔEN is 0.5 to anything LESS THAN 1.7

Question 15

Non-Polar Covalent Bond

Answer 15

-electrons are shared EQUALLY between the atoms
-ΔEN is anything LESS THAN 0.5

Question 16

Electronegativity

Answer 16

defines how strongly an atom attracts a pair of electrons it shares with another atom in a covalent bond

Question 17

Criteria for entire molecule to be polar

Answer 17

More than one polar bonds and they DO NOT cancel each other out

Question 18

INTRAmolecular Forces (definition, characteristics)

Answer 18

-forces WITHIN a molecule (ie bonds)
-can be formed or broken by CHEMICAL changes
-influence CHEMICAL properties
-relatively STRONG

Question 19

INTERmolecular Forces (definition, characteristics)

Answer 19

-forces BETWEEN molecules
-can be formed or broken by PHYSICAL changes
-influence PHYSICAL properties (ie melting point, state, etc)
-relatively WEAK

Question 20

Dipole-Dipole Forces (definition, strength, traits)

Answer 20

Caused by: dipoles of POLAR molecules that position their positive and negative ends near each other
Experienced by: Polar molecules
Strength depends on: distance between dipoles

Increase as bond polarity increases (compare ΔEN)

Question 21

Hydrogen Bonds (definition, strength)

Answer 21

Caused by: Hydrogen atom bonded to a highly electronegative atom (O, N, F) attracted to a partially negative atom
Experienced by: water, ammonia, hydrogen flouride
Strength depends on: polarity of bonds and closeness of molecules

Question 22

London/Dispersion Forces (definition, strength, traits)

Answer 22

Caused by: non-polar molecules electrons are in motion and produce temporary “dipole-like” arrangement of charge. Instantaneous dipoles created by movement of e-
Experienced by: All substances, only mentioned if other forces are absent
Strength depends on: Molar Mass

Increase as number of electrons increases, equal for isoelectronic molecules, NON-POLAR molecules ONLY experience LDF

Question 23

Rank of Intermolecular Forces Strength, Other Qualities (ranking, m.p, and b.p)

Answer 23

Strength Ranking: Ionic>H bonding>Dipole-Dipole>LDF

The stronger the IMF, the greater the energy required to overcome the forces, and therefore increased melting and boiling point.
“like dissolves like” also still applies

Question 24

Three Factors Determining Properties of a Substance

Answer 24

-The way particles bond together
-The forces that act within and among the compounds formed
-The shapes that result from these interactions

Question 25

Ionic Solids (definition, example, b.p, electrical conductivity, hardness)

Answer 25

Forms between metal and non-metal, ions arrange in crystal lattice structure with alternating + and - charges
Example: NaCl
Boiling Point: tends to be high
Conductivity: Only when dissolved in water
Hardness: hard brittle solid

Question 26

Metallic Solids (definition, example, b.p, electrical conductivity, hardness)

Answer 26

Solid with closely packed atoms held together by electrostatic interactions of free moving e-, do not all have the same properties
Example: Solid gold, solid Al
Boiling Point: some are room temp, some are extremely high
Conductivity: High
Hardness: some are hard, some are soft

Question 27

Molecular Solid (definition, example, b.p, electrical conductivity, hardness)

Answer 27

Solid composed of individual molecules held together by IMF of attraction
Example: water (ice)
Boiling Point: low
Conductivity: non-conductor
Hardness: little hardness

Question 28

Covalent Network Structure (definition, example, b.p, electrical conductivity, hardness)

Answer 28

Solid in which the atoms form covalent bonds in an interwoven network, different ones have different properties
Example: diamond, graphite, or SiO2
Boiling Points: Very high
Conductivity: not good conductors
Hardness: very hard

Question 29

Electron Sea Theory of Metallic Bonding

Answer 29

electrons move freely around positively charged nuclei