Atoms

Atoms, Isotopes, and Ions

4.3 – 4.7, 4.10, 4.11

Atomic Theory

• In 1808 John Dalton proposed atomic theory.

• Dalton’s theory explained several laws known at the time.– Law of conservation of matter– Law of definite proportions– Law of multiple proportions

Dalton’s Atomic Theory (1808)

1. Elements are made of tiny particles called atoms.

2. Atoms of a given element are identical.3. Atoms of different elements differ from

each other in some fundamental way.


4. Atoms of one element can join with atoms of other elements to form compounds.

• A given compound is always made of the same elements combined in the same ways. Explains the law of multiple proportions and the

law of definite composition.


5. Atoms are indivisible in chemical reactions.

• Chemical reactions change how atoms are grouped (bonded) together. Explains the law of conservation of matter.

Atomic Theory

• Dalton’s proposal lead to much research as to the nature of the atom.

• In the late 1800’s chemists/physicists determined that the atom is made up of smaller, subatomic, particles.

Atomic Theory - 1910

• ~1896, JJ Thomson demonstrated that atoms can emit negative particles.– Called these particles electrons.– Since atoms are neutral he also proposed that

they must contain positive particles.• These + particles were not fully described/named

until 1919.


• ~1910 Lord Kelvin proposed the “plum pudding” model of the atom.– Proposed that electrons were scattered within

a “cloud”/pudding of positive charge.


• ~1911 an experiment was conducted in Ernest Rutherford’s lab that showed the “plum pudding” model to be incorrect.– Experiment was conducted by Geiger and

Marsden and the findings interpreted by Rutherford.

– See page 84

The gold foil experiment

• What they did – see board• What they found – see board• What Rutherford concluded.

Rutherford’s Model of the Atom

• First to propose a nuclear atom.• Rutherford proposed that:

– the atom must have nearly all its mass, and positive charge, in a central nucleus about 10,000 times smaller than the atom itself.

– Most of the atom is empty space and the electrons are scattered through out this empty space.

A New Model of the Atom

Expected based onPlum pudding model

Rutherford’s modelBased on ”his” results

http://en.wikipedia.org/wiki/Image:Rutherford_gold_foil_experiment_results.svg

Subatomic Particles

• Rutherford continued to study the atom and the positive matter of the atom.– 1919, + particle named the proton

• ~1932 James Chadwick proposed the existence of a third subatomic particle, the neutron.

Subatomic Particles

Subatomic Particle

Charge Mass, amu Location in atom

Electron (e-)

-1 0 amu Outside of nucleus

Proton (p) +1 ~1 amu Nucleus

Neutron (n) 0 ~1 amu Nucleus

Mass of Subatomic Particles

• Protons and neutrons have ~ the same mass (in the range of 10 -24 g).

– Neutrons are slightly heavier.– Mass is expressed in amu

• Atomic mass unit (amu) – 1/12 the mass of a carbon-12 atom

Mass of Subatomic Particles

• The mass of the electron is tiny as compared to that of the proton and neutron.– Therefore, the electron’s mass is considered

to be ~0 amu when calculating the mass of an atom.

Subatomic Particles and the Elements

• Each element has a unique number of protons.– Number of protons defines the element.

Subatomic Particles and the Elements

• Since atoms are neutral, for every proton there is a/n _________.

• When atoms interact to form compounds, it is their ___________ that “intermingle”.

Terms

• Atomic number = number of protons in an atom– Also indicates the number of electrons in the

atom.– Finding atomic number on the periodic table.

Terms

• Mass number = sum of the # of protons and the # neutrons in the nucleus of an atom– FOR MOST ELEMENTS THE MASS

NUMBER IF NOT ON THE PERIODIC TABLE.

• You will be given enough information to determine mass number or number of neutrons.

Terms

• Isotopes = atoms of a given element that differ in mass number– Isotopes have the same number of

_____________.– Isotopes differ in the number of _______.

Isotopes

• Writing atomic symbols for isotopes– See board and pg 87

FAQ - Isotopes

• When is mass number found on the periodic table?

• What’s the atomic mass? Is it the same as the mass number?

Practice

• Start # 42 on page 110.

Ion Formation

• Ions are formed when atoms gain or lose electrons.– Proton and neutron number are unchanged

when an ion forms.

Ions - Terms

• Ion – charged atom or group of atoms– Cation = positively charged ion

• Metals form cations.

– Anion = negatively charged ion• Nonmetals form anions.

Ions

• Na atom_____ protons _____ electrons

• Na+ ion _____ protons _____ electrons

Name of ion: sodium ion

Ions

• Calcium atom_____ protons _____ electrons

• Ca 2+ ion _____ protons _____ electrons

Name of ion: calcium ion

Ions

• Sulfur atom_____ protons _____ electrons

• S2- ion _____ protons _____ electrons

Name of ion: sulfide ion

Ion Charge and the Periodic Table

Group # Ion Charge # e- lost/gained

I A 1+ 1 e- lost

II A 2+ 2 e- lost

III A metals 3+ 3 e- lost

V A nonmetals 3- 3 e- gained

VI A nonmetals 2- 2 e- gained

VII A 1- 1 e- gained

Naming Ions

• Name of a monatomic cation is the name of the element– Examples:

• Ca 2+ calcium ion• Al 3+ aluminum ion• K+

Naming Ions

• Monatomic anions are named by changing end of the name of the element to “ide”

Example: S2- sulfide ion

Naming Ions

• You need to know:N3- nitride ionP3- phosphide ionO2- oxide ionS2- sulfide ionF- fluoride ionCl - chloride ionBr- bromide ionI- iodide ion

Ionic Compounds

• Structure– In an ionic compound there is a regular

arrangement of oppositely charged particles. – Ions are arranged in a 3-D crystalline

structure that maximizes attractive forces and minimizes repulsive forces.

• Also called a lattice structure• See page 102

Ionic Compounds

• Physical Properties – all are related to the structure of the compounds– Solids at room temperature– Relatively high melting and boiling points– No vapor pressure

• Meaning… they don’t evaporate– Electrolytes

• Conduct electricity when melted or dissolved in water

Ionic Compounds

• The chemical formula for an ionic compound represents the lowest, whole number ratio of the component ions that has a net charge of zero.Total positive charge = total negative charge

Ionic Compounds

• Name the compound by naming the ions.

Ionic Compounds

• Writing formulas for and naming binary ionic compounds– Magnesium oxide

Ionic Compounds

Magnesium oxide– The formula is the simplest ratio of ions that

have a net charge of zero.– Ions present: Mg2+ and O2-

– Formula:

Ionic Compounds

Magnesium chloride– The formula is the simplest ratio of ions that

have a net charge of zero.– Ions present: Mg2+ and _____

– Formula:

Ionic Compounds

• Practice– Note we are currently applying the content of

4.11 and 5.2 (type I binary ionic compounds)

Types I Binary Compounds

• Compound between a metal and a nonmetal– Metal forms only one ion

• Name the cation and then the anion.– Name of the cation is the name of the element– Name of the anion is the name of the

nonmetal with the ending changed to “ide”

Monoatomic cations to know

Group # Charge on ion examples

IA +1 Na1+ sodium (ion)K1+ potassium (ion)

IIA +2 Mg2+ magnesium (ion)

IIIA metals

+3 Al3+ aluminum (ion)

Monoatomic anions to knowGroup # Charge on ion examples

VA -3 N3- nitride (ion)P3- phosphide (ion)

VIA -2 O2- oxide (ion)S2- sulfide

VIIA -1 F1- fluoride (ion)Cl1- chloride (ion)Br1- bromide (ion)I1- iodide (ion)

Practice

• Name chemical formula

• Chemical formula name

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Atoms