Electronic Structure and Periodic Table

Question 1

Orbital structure of hydrogen atom

Answer 1

The hydrogen electron orbits the nucleus, and the electron exists in a spherical probability cloud around the nucleus.

Question 2

Principal quantum number n

Answer 2

The number “n” defines what shell the electron is in. The higher “n” shells indicate higher energy.

Question 3

Number of electrons per orbital

Answer 3

There are n^2 squared orbitals per shell, and 2 electrons per orbital, so 2n^2 electrons per shell.

Question 4

Ground state

Answer 4

Electrons are normally in their ground state. Excited electrons come down to ground state via release of energy

Question 5

Excited state

Answer 5

When electrons in their ground state absorb energy, they get promoted to excited states

Question 6

Absorption spectra

Answer 6

The absorption spectrum shows what wavelengths of light are absorbed, and looks like black lines on a rainbow background. They correspond to the emission spectrum in pattern.

Question 7

Emission spectra

Answer 7

The emission spectrum shows what wavelengths of light are emitted. They look like colored lines on a black background. The spectrum shifts to a slightly longer wavelength.

Question 8

Quantum number l

Answer 8

The number “l” is the angular momentum for quantum numbers ranging from 0 to n-1. “l” tells us whether it’s in the s (l=0) subshell, p (l=1) subshell, d (l=2), subshell, or f (l=3) subshell.

Question 9

How many electrons exist in each subshell

Answer 9

The s subshell holds 1 orbital, the p subshell holds 3 orbitals, the d subshell holds 5 orbitals, and the f subshell holds 7 orbitals. Each hold a maximum of 2 electrons per orbital.

Question 10

Quantum number m

Answer 10

The number “m” is the magnetic quantum number that range from “-l” to “l”, including zero.

Question 11

Quantum number s

Answer 11

The number “s” is the quantum spin number, which is either +0.5 or -0.5

Question 12

Common names and geometric shapes for orbitals s, p, d

Answer 12

For “s”, this equates to one orbital, “l” = 0, and a spherical shape. For “p”, this equates to 3 orbitals, “l” = 1, and a fission shape. For “d”, this equates to 5 orbitals, “l” = 2, and a clover-like shape.

Question 13

Arrangement of subshells

Answer 13

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

Question 14

Orbital diagrams

Answer 14

Each subshell gets a row where each +0.5 spin and each -0.5 spin are an arrow.

Question 15

Aufbau principle

Answer 15

Shells and subshells of lower energy get filled first.

Question 16

Hund’s rule

Answer 16

When you fill a subshell with more than 1 orbital, you must fill each orbital with a single electron with the same spin.

Question 17

Pauli Exclusion Principle

Answer 17

2 electrons in the same orbital must be of different spins

Question 18

Bohr atom

Answer 18

Electron orbiting nucleus in a circular orbit and the larger “n” values have a larger orbiting radii.

Question 19

Effective nuclear charge

Answer 19

The nuclear charge minus the shielding charge. The higher the efective nuclear charge, the more stable it is. It increases for outer electrons as you go left to right in the periodc table. Shielding electrons are those that stand between the nucleus and the electron. They are lower in energy.

Question 20

Alkali metals

Answer 20

Group 1. Single valence electron, low ionization energy, very reactive, more reactive as you go down because of increasing radii. Most commonly found in +1 oxidation state

Question 21

Alkaline earth metals

Answer 21

Group 2. Two valence electrons, low ionization energy, quite reactive, more reactive as you go down, and most commonly found in the +2 oxidation state

Question 22

Halogens

Answer 22

Group 7. Seven valence electrons (2 s subshell and 5 p subshell). High electron affinity. Very reactive. More reactive as you go up because of decreasing radii. Most commonly found in -1 oxidation state.

Question 23

Noble gases

Answer 23

Group 8. Eight valence electrons. High ionization energy and low electron affinity. Inert. Have oxidation state of 0.

Question 24

Transition metals

Answer 24

High conductivity due to loosely bound outer d electrons. Varied oxidation states, but always positive.

Question 25

Representative elements

Answer 25

Include the s block and the p block of the periodic table. Valence shell fills left to right.

Question 26

Metals

Answer 26

Metals like to lose electrons, and have lower electronegativity. Form basic oxides. Good conductor of heat and electricity. Malleable, ductile, luster, and solid.

Question 27

Nonmetals

Answer 27

Likes to gain electrons to form a negative oxidation state. Higher electronegativity, and forms acidic oxides. Poor conductor of heat and electricity.

Question 28

Oxygen and sulfur relationship

Answer 28

Oxygen and sulfur are chemically similar.

Question 29

First and second ionization energies

Answer 29

First ionization energy is the energy needed to knock off the first valence electron. Second ionization energy is the energy needed to knock off second valence electron.

Question 30

Prediction of first and second ionization energies from electronic structure for elements in different groups or rows

Answer 30

Decreases as you go down (increasing radii) and increases as you go right (decreasing radii). Highest are noble gases and lowest are alkali metals. Second ionization energy is ALWAYS higher than the first, although alkaline earth metals have a realtively lower secondary ionization energy.

Question 31

Electron affinity

Answer 31

Amount of energy released when something gains an electron

Question 32

Electron affinity variations with group or row

Answer 32

Decreases as you go down , and increases as you go right. Highest peaks are halogens, and lowest are noble gases.

Question 33

Electronegativity

Answer 33

How much something hordes electrons in a covalent bond. Increases toward the top right. If electronegativity is too great, ionic bonds form.

Question 34

Electron shells and sizes of atoms

Answer 34

Going down the periodic table means jumping to the next shell. Effective nuclear charge decreases because old shell stands in between nucleus and new shell. Atomic size increases going down a group. Size decreases going right.

Question 35

Diamagnetic atoms

Answer 35

Contains an even number of electrons and have all occupied subshells filled. Repelled by magnetic field.

Question 36

Paramagnetic atoms

Answer 36

Unpaired valence electrons and attracted to the magnetic field.

Question 37

Positive charge and atomic radii

Answer 37

The more positive, the smaller the atomic radii

Question 38

Hydrogen bonding occurs with…

Answer 38

Hydrogen, Oxygen, Nitrogen, and Fluorine

Question 39

Differentiate between covalent and ionic bonds

Answer 39

Ionic bonds are usually a metal and a non-metal. Covalent bonds are usually two non-metals.