Unit 1: Atomic Structure & Electron Configuration

I. Theories and Models Scientific Model – A pattern, plan,

representation or description designed to show the structure or workings of an object, system or concept.

A. Greeks • 400 B.C.• Democritus• particle theory- matter could not be divided into

smaller and smaller pieces forever, eventually the smallest possible piece would be obtained and would be indivisible.

• called nature’s basic particle atomos-indivisible• no experimental evidence to support theory

B. John Dalton• 1808• English school teacher• Established first atomic theory:

1. Matter is composed of atoms.2. Atoms of a given element are identical to each other, but different

from other elements.3. Atoms cannot be divided nor destroyed.4. Atoms of different elements combine in simple whole-number ratios

to form compounds.5. In chemical reactions, atoms are combined, separated or rearranged.

• Model: tiny, hard, solid sphere

C. JJ Thomson• 1897• cathode ray tube experiment• given credit for discovering electrons,

resulting in the electrical nature of an atom• Plum pudding model – sea of positive charges

with negative charges embedded evenly throughout.

D. Ernest Rutherford• 1911• Gold Foil (Alpha Scattering) Experiment

• Conclusions: atom is mostly empty space most of mass of atom is in the nucleus nucleus is positively charged

• Model:

E. Niels Bohr• 1913• Rutherford’s student• electrons arranged in energy levels (orbits)

around the nucleus due to variation in energies of electrons

• higher energy electrons are farther from nucleus

• Planetary Model:

F. Quantum Model• 1924-current• Collaboration of many scientists• Better than Bohr’s model because it describes

the arrangement of e- in atoms other than H• Based on the probability (95% of time) of

finding and e- or an e- pair in a 3D region around the nucleus known as an orbital

• Model (on board)

II. General Structure of Atom

nucleus center of atom p+ & n0 located here positive charge most of mass of atom, tiny

volume very dense

e- cloud surrounds nucleus e- located here negative charge most of volume of atom,

negligible mass low density

III. Quantification of the AtomA. Atomic Number - the number of p+ in nucleus

All atoms of the same element have the same atomic number.

Periodic table is arranged by increasing atomic number.

if atom is electrically neutral, then the

#p+ = #e-

B. Mass Number - the total number of p+ & n0 in nucleus of an atom.

Round the atomic weight to a whole number n0 = mass number - atomic number

C. Ions – atoms of an element with the same number of p+ that have gained or lost e-, therefore having a – or + charge

atoms form ions in order to be more stable like the noble gases anion – ion with negative charge (gained e-)

• non-metal elements tend to form anions (ex. S2-)• change the end of the element name to –ide (sulfide ion)

cation – ion with a positive charge (lost e-) • metal elements & H tend to form cations (ex. Sr2+)• Roman numerals may be used in the name of some metal ions that

can lose various numbers of e- (ex. Tin (IV) ion)

D. Isotopes – atoms of an element having the same number of p+, but a different number of n0, resulting in a different mass number.

• Two ways to represent isotope symbols:

or C-14

• Write mass # after the element name:

carbon-14

C146

mass #

atomic #mass #

Isotopes of Hydrogen

Name Symbol e- n0 p+ Mass # Atomic #

Hydrogen-1 (protium) 1 0 1 1 1

Hydrogen-2 (deuterium) 1 1 1 2 1

Hydrogen-3 (tritium) 1 2 1 3 1

H11

H21

H31

E. Average Atomic Mass – weighted average of all natural isotopes of an element expressed in amu* (atomic mass units).

based on % abundance of isotopes steps for calculating:

1. change % to decimal

2. multiply decimal and mass number

3. add all results

4. place amu unit with answer

*amu=1/12 mass of C-12 isotope

IV. Electromagnetic RadiationA. Properties

1. Form of energy which requires no substrate to travel through.

2. Exhibits properties of a sine wave

amplitude wavelength (λ)crest

trough

line of origin

a. wavelength = distance between consecutive crests (Greek letter lambda = λ)

b. frequency = # wave cycles passing a given point over time (seconds); (Greek

letter nu = ν )

*measured in Hertz (Hz)= 1/s, s-1, or per second

c. all types of ER travel in a vacuum at the speed of light (c) = 3.00 x 108 m/s

3. light equation

c=λν* λ & ν are inversely (indirectly)

proportional (as one increases, the other decreases)

λ

ν

*energy & ν are directly related (as one increases/decreases, so does the other

*energy equation: E=hν

h= Plank’s constant = 6.63 x 10-34 J·s

V. Emission/Absorption Spectra*The e- is the only SAP that absorbs/emits

energy.A. Absorption Spectrum –when an e- absorbs

energy, it moves from the ground state (most stable arrangement of e-) to an excited state (which is not stable)

B. Emission Spectrum - when an e- emits energy, it falls from the excited state back to ground

state, releasing energy in the form of electromagnetic radiation, which may be visible

*unique to each atom http://chemistry.bd.psu.edu/jircitano/periodic4.html

VI. Electron ConfigurationA. Describes the arrangement of e- in an atom

1. each main energy level is divided into sublevels

2. each sublevel is made up of orbitals, each of which can hold up to 2 e-

*chart

Sublevel # of orbitals

shape

s 1

p 3

d 5

f 7

3. due to main energy levels getting closer together, sublevels overlap

4. Aufbau principle – states that e- fill orbitals of lower energy sublevels

first

5. Abbreviated Configurations – use the preceding noble gas symbol (in brackets) to represent the filled inner core of e-. Then write the remaining configuration for the atom.

6. Orbital Configurations- arrangement of e- within sublevels

2 rules determine arrangement: a. Hund’s Rule – each orbital within a sublevel receives 1 e- before it gets 2

* orbitals in the same energy sublevel are degenerate (of equal energy)

b. Pauli Exclusion Principle – no 2 e- in an orbital can have the same spin.

= clockwise spin = counterclockwise spin

*exceptions