Chapter One: Atomic Theory

Question 1

John Dalton’s View of the Atom

Answer 1

The atom was a billiard ball-shaped structure. No charges or sub-atomic particles

Question 2

J.J. Thomson’s View of the Atom

Answer 2

• “Plum pudding Model”

- atom was a positively charged molecule with negative ions randomly assorted throughout

Question 3

J.J. Thomson’s Experiment

Answer 3

• Cathode Ray Tube
• Created a tube with a positively charged anode on one side and a negatively charged cathode on the other
• Applied a magnet to the middle of the tube and the beam was displaced
• Negatively charged particles were emanating to the positive field

Question 4

Ernest Rutherford’s View of the Atom

Answer 4

The atom was mostly empty space, with electrons surrounding a dense, positively charged nucleus. The electrons were no longer randomly assorted

Question 5

Ernest Rutherford’s Experiment

Answer 5

Gold Foil Experiment

• emitted alpha particles towards a thin gold sheet (chosen for its density). Put a deflecting screen around the foil to detect the path of the electrons
• most alpha particles went straight through but some deflected backwards or sideways

Question 6

The Electromagnetic Spectrum

Answer 6

Energy that has specific combinations of energy and wavelength
Specific colour = Specific amount of energy

Question 7

The Electromagnetic Spectrum Rules (2)

Answer 7

1. Only certain energies are possible

2. Brightness depends on the number of photons emitted

Question 8

Niels Bohr’s View of the Atom

Answer 8

-electrons orbit the nucleus in circular, predictable orbits
-the larger the orbit, the more energy it possessed
-

Question 9

According to the Bohr hydrogen model, what happens when an electron receives energy?

Answer 9

When an electron receives additional energy, its moves from its original state to the next energy level. To return to its home state it releases energy (a quanta), which gives off a specific colour

Question 10

Louis de Broglie

Answer 10

• Theorized that if light can be both a wave and a particle, electrons can too.
• if an electron has wavelike motion and a fixed radius, only certain frequencies, wavelengths, and energies are possible

Question 11

Werner Heisenberg’s Uncertainty Principle

Answer 11

The location of the electron can’t be precisely determined. Location began to be measured in PROBABILITY

Question 12

Main Principles of the Quantum Mechanical Model (4)

Answer 12

1. Electrons occupy the space surrounding the nucleus and exist in several discreet principal energy levels
2. Electrons in higher principal levels have more energy
3. Each principal energy level consists of energy sublevels with slightly different energy values (orbitals)
4. Orbitals are regions of space with higher probabilities of electrons (2 e-/orbital)

Question 13

Principal Energy Level

Answer 13

Main orbital level, represented by the quantum number n

Question 14

Three Characteristics that Define Sub-Orbitals

Answer 14

1. Size (or amount of enrgy)
2. Shape (sphere, dumbbell, etc)
3. Spatial Orientation (x-y-z axis)

Question 15

The Aufbau Principle

Answer 15

Each electron occupies the lowest energy orbital possible

Question 16

The Pauli Exclusion Principle

Answer 16

A maximum of two electrons may occupy a single electron orbital, if they have opposite spins

Question 17

Hund’s Rule

Answer 17

Each orbital must hold one electron before any orbital holds a second, if they have opposite spins

Question 18

The Periodic Law

Answer 18

When the elements are arranged by atomic number, the physical and chemical properties vary periodically.

Question 19

Atomic Radius

Answer 19

Size of the atom; measured by taking half the distance between two nuclei of atoms in a solid crystal.

Question 20

Periodic Table Trends for Atomic Radius

a. Moving Down a Group
b. Moving Across a Row/Period

Answer 20

a. Atomic radii increases as you move down a group, due to the shield of lower electrons between the nucleus and the valence electrons
b. Atomic radii decreases as you move across a group, since each atom is added to the same principal energy level with no extra shield

Question 21

Ionization Energy

Answer 21

The amount of energy is takes to detach one electron from a gaseous atom/ion

Question 22

Periodic Table Trends for Ionization Energy

a. Moving Down a Group
b. Moving Across a Row/Period

Answer 22

a. IE decreases as you move down a group, since the shield of lower shells makes it easier to detach electrons
b. IE increases as you move across a period, since there is greater electrostatic attraction

Question 23

Electrostatic Attraction

Answer 23

The force between the positive nucleus and the negative valence electrons.

Question 24

Electrostatic Repulsion

Answer 24

The force between valence electrons in the outermost shell.

Question 25

Electronegativity

Answer 25

The ability of an atom to attract bonding electrons to itself. Inverse of IE.

Question 26

Covalent Bonds (+ their EN range)

Answer 26

Electrons are shared equally.
EN low (approx 0)

Question 27

Polar Covalent Bonds (+ their EN range)

Answer 27

Electrons are shared unequally with partial charges on each atom.
EN medium (1-1.6)

Question 28

Ionic Bonds (+ their EN range)

Answer 28

Electron transfer.
EN high (greater than 2)