Chapter 4

Chapter 4AtomsSection 1: The Development of Atomic Theorypp. 113 - 118Beginnings of Atomic Theory

First theory that was proposed 2,000 years ago

b. Democritus suggested that the universe was made of indivisible units he called atoms.

c. He did not have evidence to convince people of his theory

B. Daltons Atomic Theory

a. Proposed in 1808 by John Dalton (English schoolteacher)

b. All atoms of a given element were exactly alike, and atoms of different elements could join to form compounds

C. Thomsons Model of the Atom

a. Developed in 1897 by J.J. Thomson (British scientist)b. Cathode rays: mysterious rays in vacuum tubesc. Thomson was studying these cathode rays and it was discovered that atoms were not indivisibled. Cathode ray tube experiment suggested that cathode rays were made of negatively charged particles that came from inside atomsi. Atoms could be divided into smaller parts!

Cathode Ray ExperimentD. Rutherfords Model of the Atoma. Ernest Rutherford (British scientist)

b. Proposed that most of the mass of the atom was concentrated at the atoms center

c. Atoms positive charge is concentrated in the center of the atom (nucleus) and negative electrons orbitRutherford's Experiment

Section 2: The Structure of the Atompp. 119 - 127What is in an atom??

a. Three main subatomic particles are distinguished by:- mass- charge- location in the atom

b. Nucleus: center of each atom, small and dense, made of protons and neutronsc. Protons: positive charged. Neutrons: no chargee. Electrons: move around outside the nucleus in a cloudf. Mass of electron is much smaller than that of a proton or neutron

ParticleChargeMass (kg)Location in atomProton+11.67 x 10 -27In the nucleusNeutron01.67 x 10-27In the nucleusElectron-19.11 x 10-31Outside nucleus

g. Each element has a unique number of protons

i. elements are defined by the number of protons in an atom of that element

h. Most atoms have an equal number of protons and electrons, whose charges cancel out

Example: He (helium) atom

Charge of two protons: +2Charge of two neutrons: 0Charge of two electrons: -2_ 0ii. If an atom gains or loses electrons, it becomes charged (ion)

Electric force holds atom together

i. Electric force: positive and negative charges that attract each other

B. Atomic Number and Mass Number

a. Atoms of each element have the same number of protons, but they can have different numbers of neutrons

b. Atomic Number (Z): number of protons

i. Each element has a unique atomic number

ii, hydrogen has only one proton, so Z=1iii. Uranium has 92 protons, Z = 92

iv. Atomic number of an atom NEVER changes

c. Mass number (A): number of protons plus the number of neutronsi. Fluorine has 9 protons and 10 neutrons, so A=19ii. Oxygen has 8 protons and 8 neutrons, A=16

iv. Although atoms of an element have the same atomic number, they can have different mass numbers because the number of neutrons can vary.

What is Lithiums A?What is Lithiums Z?

C. Isotopes

a. Isotope: atom that has the same number of protons but a different number of neutrons relative to other atoms of the same element

b. have different masses because numbers of neutrons differs

c. Hydrogen isotopes:

i. Protium: A=1, Z=1ii.Deuterium: A=2, Z =1iii. Tritium: A=3, Z=1

Protiumd. Some isotopes are more common than othersi. most common isotope of hydrogen is protium

Protiume. Radioisotopes: emit radiation and decay into other isotopes

i. Tritium is an unstable isotope of hydrogen and decays over time

ii. Continue to decay until the isotope reaches a stable point

f. Number of neutrons can be calculated

i. write the mass number and atomic number of the isotope before the symbol of the element

Protiumii. Number of neutrons can be found by subtracting the atomic number from the mass number

iv. Example: uranium (U) -235

Mass number (A) = 235 - Atomic number (Z) = 92_ 143

D. Atomic Massesa. because working with tiny masses is difficult, atomic masses are usually expressed in unified mass units (u) or atomic mass unit, amu

b. Example: Carbon-12i. Isotope of carbonii. Has 6 protons and 6 neutronsiii. Each individual proton and neutron has a mass of about 1.0 u

Section 3: Modern Atomic Theorypp. 128 - 132Modern Models of the Atom

a. In 1913, Bohr (Danish physicist), proposed that the energy of each electron was related to the electrons path around the nucleus

i. Can only be in certain energy levels

ii. Must gain energy to move to a higher energy level

iii. Must lose energy to move to a lower energy level

b. Exact location of an electron cannot be determinedi. Orbital: shaded region where there is a likelihood of finding an electron, the darker the shading, the better the chance of finding an electron

B. Electron Energy Levels

a. The number of energy levels that are filled in an atom depends on the number of electrons

b. Valence electrons: electrons in the outer energy level of an atom

i. Determine the chemical properties of an atom

c. Four types of orbitalsi. s, p, d, and f

ii. s orbitals are the simplest

1. only one possible orientation in space

2. shaped like a sphere

3. lowest energy

4. can hold up to 2 electrons

iii. p orbital1. shaped like a dumbbell

2. can be oriented in space in 1 of 3 ways

3. each p orbital can hold up to 2 electrons for a total of 6 electrons

iv. d and f orbitals

1. more complex

2. 5 possible d orbitals

3. 7 possible f orbitals

4. each orbital can hold a maximum of 2 electronsd. Orbitals determine the total number of electrons that each level can holdOrbitals and Electrons for Energy Levels 1-4Energy LevelNumber of Orbitals by TypeTotal Number of Orbitalsspdf111 = 12131 + 3 = 431351+3+5 = 9413571+3+5+7 = 16Number of Electrons261832C. Electron Transitions

a. electrons jump between energy levels when an atom gains or loses energyb. Ground State: lowest state of energy of an electron

c. Excited State: when an electron gains energy

d. Gain energy by absorbing a particle of light (photon)