Unit 3: Atomic Structure & The Periodic Table. Chapters 13 & 14. Energy: The capacity to do work. Potential Energy: stored energy due to position or condition. Chemicals can store energy; thus they have potential energy. Kinetic Energy: energy in motion. Kinetic Theory.
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13 & 14
1. All elements are composed of tiny indivisible particles called atoms.
2. Atoms of the same element are identical. The atoms of any one element are different from those of any other element.
3. Atoms of different elements can physically mix together or can chemically combine.
4. Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of element are never changed into atoms of another element as a result of a chemical reaction.
J.J. Thomson:revised Dalton’s model by proposing that electrons were stuck to the outside of the atom.
Structure of Atoms:
-(Discovered by E. Goldstein using canal rays. Canal rays traveled from the positive metal plate to the negative metal plates)
- Discovered by James Chadwick
- Discovered by J.J. Thomson using a cathode ray - the rays were attracted to a metal plate of positive charge.
- the # of protons = the # of electrons
- To find the # of neutrons subtract the mass # from the atomic #
Mass number = 14
Atomic number = 7
# of protons = 7
# of electrons = 7
# of neutrons = 14 – 7= 7
284Uus (283.4 amu) 34.60%
285Uus (284.7 amu) 21.20%
288Uus (287.8 amu) 44.20%
What would the average atomic mass be?
Examples on the board
1. Principal Quantum Numbers ( ) = 1,2,3,4….
2. Each principal level contains sublevels
* Table 13.1 p. 364
3. Atomic orbitals are regions where electrons can be found. (Letter denotes the orbital)
( exist in three different planes)
a. N = 1; 1 sublevel; 1s orbital
b. N = 2; 2 sublevels; 2s (1 orbital), 2p ( 3 orbitals)
c. N = 3; 3 sublevels; 3s ( 1 orbital), 3p (3 orbitals), 3d (5 orbitals)
d. N = 4; 4 sublevels; 4s (1 orbital), 4p (3 orbitals), 4d ( 5 orbitals), 4f (7 orbitals)
(n = principal quantum number).
This equals the maximum # of electrons that the sublevel can hold.
1. Unstable systems tend to lose energy to become stable.
2. Electrons try to form stable arrangements with the nucleus.
3. The way in which electrons are arranged around the nuclei of atoms is called electron configuration.
a. Aufbau principle: electrons enter orbitals of lowest energy level first.
b. Pauli exclusion principle: and atomic orbital may describe at most two electrons. (arrows show the direction of electron spin)
c. Hund’s rule: when electrons occupy orbitals of equal energy, one electron enters each orbital, all of orbitals contain one electron with parallel spins.
5. When writing electron configurations for ions you must add or subtract the # of electrons gained or lost to create the ion.
- Alkali Metals – Group 1A
- Alkaline Earth Metals – Group 2A
Transition Metals & Inner Transition Metals – make up Group B (1B – 8B)
- Halogens – 7A
- Noble Gases – 0
1. You can determine the charge of an ion by what group it is in.
1A = +1 5A = -3
2A = +2 6A = -2
3A = +3 7A = -1
4A = +/- 4
- nuclear charge
- distance from the nucleus
Ionization energy increases across the period ( left – right) due to increased nuclear charge
- size increases down a group (top – bottom)
- size decreases across the period ( left – right)
This is due to an increase in nuclear charge pulling them closer… the energy level stays the same
How likely/vigorously an atom is to react with other substances
The farther left and down you go the easier it is for electrons to be taken away. (Higher Reactivity)
The farther right and up you go the higher electronegativity – vigorous exchange of electrons
Elements can be classified into 4 groups based on electrons.
1. Noble gases: outermost s & p sublevels are filled. Belong to group 0. (Also called inert gases.)
2. Representative elements: outermost s or p sublevel is partially filled
4. Inner transition metals: metallic elements in which the outermost s sublevel and nearby f sublevel generally contain electrons. (Lanthanide & Actinide series)
(units = hertz Hz)
Example: Calculate the wavelength of the yellow light emitted by a sodium lamp if the frequency of the radiation is 5.10 x 10 14 Hz (5.10 x 10 14 s-1).
c = 3.00 x 108 m/s
Frequency (ν) = 5.10 x 1014 s-1
wavelength (λ) = ??? m
E = h x ν
Example: Calculate the energy (J) of a quantum of radiant energy (the energy of a photon) with a frequency of 5.00 x 1015 s-1.
ν = 5.00 x 1015 s-1
h = 6.63 x 10 -34 J x s
Energy(E) = ??? J