Chapter 4 Atoms

Section 1

Development of Atomic Theory

BR. Who came up with the first theory of atoms?

• Objectives- • 1. Give an example of how new scientific

data can cause an existing scientific explanation to be supported, rejected or revised

• 2. Evaluate selected theories based on supporting scientific evidence.

Objectives (cont.)

• 3. Cite evidence that scientific investigations are conducted for many reasons.

• 4. Identify scientific evidence that has caused modifications in previously accepted theories.

• GLE’s:

Democritus

• Over 2000 years ago• Democritus• Universe made of indivisible units called

atoms.• Atomos- unable to be cut or divided• Did not have evidence to support theory

Dalton

• 1808 John Dalton revised atomic theory• Atoms could not be divided• All atoms of a given element were exactly

alike; and atoms of different elements could join to form compounds.

• Based on experimental evidence

Dalton (cont.)

• Law of definite proportions- a chemical compound always contains the same elements in exactly the same proportions by weight or mass. (supported Dalton’s theory)

• Foundation of modern atomic theory• Could not explain all experimental

evidence

Thomson

• 1897 experiment suggested that atoms were not indivisible

• He was experimenting with electricity • studying cathode rays not atoms• Cathode ray tube experiment suggested

that cathode rays were made of negatively charged particles that came from inside the atoms

Thomson (cont.)

• Revealed that atoms could be divided• Discovered electrons, negative particles• New model- electrons spread through-out

the atom ( plum pudding model)

Mass and positive charge evenly distributed

Electrons scattered through out

Rutherford

• Found Thomson’s model needing revising• Proposed that most of the mass of the

atom was in the center• Conducted Gold-foil experiment where

most particles passed straight through• Some particles were deflected• Some particles came straight back• Not what he expected

Rutherford (cont.)

• Discovered the nucleus• Nucleus was very small• Electrons orbit the nucleus (like sun and

planets)• Led to new model of the atom

Section 2- Structure of Atom

• BR. Rutherford’s Gold foil experiment led to the discovery of what?

• Objectives:• 1. Identify the 3 subatomic particles by

location, charge, and relative mass • 2. Describe the results of loss or gain of

electrons on charges of atoms

Objectives (cont.)

• 3. Identify valence electrons in first 20 elements.

GLE’S:

Atom

• Three subatomic particles compared by mass, charge, and location in the atom.

• Copy chart on page 119

• Nucleus- small, dense, center of atom• Atoms are Neutral.

Atom (cont.)

• Nucleus is made of

1. protons- + charged particle

2. neutron- neutral or no charge particle

Protons and neutrons are almost equal in size and mass.

Electrons• Move in a dense cloud (fan blades

moving) outside the nucleus• Very tiny--- 1837 electrons = 1 neutron or

proton• Negatively charged• Exact location cannot be determined.

Speed and direction cannot be determined• Located by shading; shaded region is

orbital: darker shading better chance to find

Protons

• Each element has a unique number of protons

• Elements are identified by the number of protons that they have in the nucleus of an atom

• Atoms are neutral because they have the same number of p and e. They cancel each other out.

Ions

• Atoms that have lost or gained electrons.

• Lose electrons become positive.

• Gain electrons become negative

• If atoms lose or gain electrons they are not atoms, but IONS.

Atoms

• Atoms are held together by an electric force

+ and – charges attract each other by an electric force

This attraction is what holds the atom together just like the attractive force between solids and liquids.

Atoms and Elements• Atoms of different elements have unique

structures.

• Because atoms have different structures, they have different properties.

• Atoms of the same element can vary in structure also.

• Atoms of each element have the same number of protons, but different numbers of neutrons.

Atomic number• Atomic number equals the number of

protons.• Since atoms are neutral, it also equals the

number of electrons.

• Neutral atom-- + = -

Atomic numbers go from 1 to 116

Mass Number

• Equals the total number of subatomic particles in the NUCLEUS of the atom.

• Nucleus contains p and n.• Mass number is equal to p + n.• Neutrons vary so mass number can vary

for the same element.

• See figure 3 on page 121

Isotopes

• Isotope has same atomic number but different number of neutrons.

• Isotopes have same atomic number but different mass numbers.

• Isotopes have the same number of protons but different number of neutrons

Isotopes (cont.)

• Some isotopes are more common than others.

• Radioisotopes- unstable isotopes that emit radiation and decay into other isotopes.They continue to decay until they reach a

stable isotope.They decay at a fixed rate. (fraction of a

second to millions of years)

Isotopes (cont.)

• Since isotopes have the same number of protons and electrons, they have the same chemical properties.

• Isotopes have different masses.

• Isotopes of an element vary in mass because their numbers of neutrons differ.

Isotopes of Water

• Water has 3 isotopes. Each isotope has 1 proton and 1 electron.

• Protium- 1 proton and no neutrons (mass number of 1

• Deuterium-1 proton and 1 neutron (mass number of 2)

• Tritium- 1 proton and 2 neutrons mass number of 3)

Calculating Neutrons

• Isotopes are written

35 mass number ( p+n)• Cl symbol

17 atomic number (p)

Mass number – atomic number = neutrons

Atomic mass

• Atoms are expressed in unified atomic mass units because the mass is so small.

• Unified atomic mass = equal to 1/12th of the mass of a carbon-12 atom.

• Also called atomic mass unit

Atomic Mass

• Average atomic mass for an element is a weighted average.

• More common isotopes have more effect

than less common isotopes of the element.

Mole

• Mole- collection of a very large number of particles.

602,213,670,000,000,000,000,000 particles

Written 6.02 x 1023 particles and is called Avogadro’s number

Avogadro’s number= the number of atoms in 12 grams of carbon-12. (Popcorn kernels covering

the US 310 miles tall)

Molar mass

• Molar mass- the mass in grams of 1 mole of a substance

1 mole C12 = 12 g

Converting between Moles and Grams

Amount x molar mass of element = Mass(g Moles 1 mole of element

3 moles x 32.07 g S = 96.21 g S S 1 mole S

96.27 g S x 1 mole S = 3 mol S 32.21 g S

Work problem 1 a-d on page 126

Compounds have Molar mass

• Add all molar masses in compounds and then work the same way.

H2O x (1.01 g H x 2) + 16 g O = 18.01 g/mol

1 mole H2O H2O

1 mol H20 = 18.02 g

Section 3Modern Atomic Theory

• BR List the 3 subatomic particles and give their location, relative mass, and charge.

• Objectives:• Describe the results of the loss or gain of

electrons on the charges of atoms.• Identify valence electrons in first 20 elements.• Draw Bohr models of 1st 20 elements• GLE’S:

Modern Model of Atom

• Electrons are found only in certain energy levels. NOT between levels

• Location of electrons can not be predicted precisely

• Bohr- electrons can be in only certain energy levels.

• Bohr energy level related to electron’s path around the nucleus

Modern Model of Atom (cont.)

• Electrons must gain energy to move to a higher energy level.

• Electrons must lose energy to move to a lower energy level.

• 1925 Bohr’s model revised.

Modern Atomic Theory (cont.)

• Old out – not like sun and planets• New in – • Electrons behave more like waves on a

vibrating string than like particles.

Energy levels and Electrons

• Many energy levels for electron to occupy• The number of energy levels that are filled

in an atom depends on the number of electrons.

• Valence electrons- those electrons in outer most energy level

• Valence electrons determine the chemical properties of the atom.

Energy levels

• Maximum electrons in energy level• 1st = 2• 2nd= 8• 3rd= 18• 4th= 32• Must fill 1st and 2nd energy level before

going to the 3rd energy level.

Energy levels (cont.)

• There are 4 types of orbitals.

• Orbitals are s, p, d, f.

• Orbitals determine the number of electrons that each level can hold.

Electron Jumping

• Electrons jump between energy levels when an atom gains or loses energy.

• Lowest energy level called ground state.

• Excited state-gains energy it moves to another level

How Electrons Move

• Electrons gain energy by absorbing a photon and move to a higher energy level

• Photon- particle of light; each have different energies

• The electron may fall back to previous energy level when it releases a photon.

• Photons determine which level the electron will• jump to.

Light-

• Photons determine which level the electron will jump to.

• Atoms absorb or emit light at certain wavelengths.

• Energy of photon is related to the wavelength of the light.

• High energy photons= short wavelengths• Low energy photons= long wavelengths

Atomic Fingerprint

• Because of each element’s unique atomic structure, the wavelengths emitted depend on the particular element.

• Each element emits its own characteristic color. Neon= red blue=copper

sodium= yellow strontium= red

orange = calcium green= barium