The Atom, The History, and The Periodic Table

(Chemical Naming and Formula Writing)

Unit #3

Part I

Chemistry

Theories Involving Matter and Atoms

• Democritus– Greek– 400 B.C. – Greeks: “all matter is

composed of 4 fundamental substances”

• Earth, air (wind), water & fire– Democritus: “matter is composed

of small, indivisible parts,” (Greek – “atomos”)

– No experiments to test; no definitive conclusion– First scientist to discover the idea of an atom

Alchemy (next 2000 years)

• Mixture of science and mysticism. • Lab procedures were developed, but alchemists did not

perform controlled experiments like true scientists.

Theories cont.

• Aristotle– Greek– Rejected the idea of atoms– Expanded on idea of 4

elements– Reasoning from logic & observation– Also in line with religion

Theories cont.

• Lavoisier- French chemist- “Father of Modern Chemistry”- Experimented and measured the masses of reactants and products of various reactions- Law of Conservation of Matter

Theories cont.

• Proust

– French chemist

– Showed that a given compound always contains the same proportion of elements by mass

– Law of Definite Proportions

Theories cont.

• John Dalton (1803)– English schoolteacher– Thought about atoms

as particles that might compose elements

– Billiard Ball Model• atom is a uniform,

solid sphere

– Elements combine in the ratio of small whole numbers

– Law of Multiple Proportions

John Dalton

Dalton’s Four Postulates

1. Elements are composed of small indivisible particles called atoms.

2. Atoms of the same element are identical. Atoms of different elements are different.

3. Atoms of different elements combine together in simple proportions to create a compound.

4. In a chemical reaction, atoms are rearranged, but not changed.

Henri Becquerel (1896)

• Discovered radioactivity– spontaneous emission of

radiation from the nucleus

• Three types:– alpha () - positive– beta () - negative– gamma () - neutral

J. J. Thomson (1903)

• Cathode Ray Tube Experiments– beam of negative particles

• Discovered Electrons– negative particles within the

atom

• Plum-pudding Model

J. J. Thomson (1903)

Plum-pudding Model– positive sphere (pudding)

with negative electrons (plums) dispersed throughout

Ernest Rutherford (1911)

• Gold Foil Experiment

• Discovered the nucleus– dense, positive charge in the

center of the atom

• Nuclear Model

Ernest Rutherford (1911)

• Nuclear Model– dense, positive nucleus surrounded by

negative electrons

Niels Bohr (1913)

• Bright-Line Spectrum– tried to explain presence of

specific colors in hydrogen’s spectrum

• Energy Levels– electrons can only exist in

specific energy states

• Planetary Model

Niels Bohr (1913)

• Planetary Model

– electrons move in circular orbits within specific energy levels

Bright-line spectrum

Erwin Schrödinger (1926)

• Quantum mechanics – electrons can only exist in

specified energy states

• Electron cloud model – orbital: region around the

nucleus where e- are likely to be found

Erwin Schrödinger (1926)

Electron Cloud Model (orbital)• dots represent probability of finding an e-

not actual electrons

James Chadwick (1932)

• Discovered neutrons– neutral particles in the

nucleus of an atom

• Joliot-Curie Experiments– based his theory on their

experimental evidence

James Chadwick (1932)

Neutron Model• revision of Rutherford’s Nuclear Model

Atoms

• Best current representation of the atom is a charged-cloud

• Smallest particle of an element that retains its properties

• Electrically neutral; # Protons = # electrons•  Parts of an atom

– Nucleus (contains both protons and neutrons)– Electron Cloud (contains electrons)

Parts of an Atom

• Nucleus– Small, dense center of positive charge. – Protons

• Positively charged particles within the nucleus

– Neutrons• Particles within the nucleus with no charge

• About the same mass as protons

Parts of an Atom cont.

• Electron Cloud– Empty Space– Holds electrons, which are

densely packed• Negatively charged particles

found outside the nucleus• Much smaller than protons

and neutrons• Proton/neutron mass = 1.67 x

10-24 g• Electron mass = 9.11 x 10-28

g

– More about electron behavior later

e-e-

e-

e-

p+

p+

p+

nonono

NucleusProtonsNeutronsElectrons

Picture of an atom

Atoms and Elements

• What is the difference between an element and an atom?

• An atom is a single example of an element.

• An element is the collective term for many atoms of a single substance.

Periodic Table of Elements Structure

• Listed in order of inc. atomic #

– Columns = Families or GROUPS

• Have similar chemical properties

• Referred to by the number and letter (A or B) over the column

• Many have special names

– Rows = PERIODS

• Contains info on physical properties (i.e. mp, bp, density, physical states, etc)

Element Characteristic • Each square usually contains

– Element name– Element symbol– Atomic number

• The number of protons• Symbolized at “Z”

– Atomic mass or mass number• Atomic mass: decimal mole weight • = to the average mass numbers of all isotopes• Mass number: rounded mole weight• Mass number: The sum of the number of the protons and the

number of neutrons– State of matter (usually)

Short Handing Element Characteristics

• SymbolizingA

zX

A = mass # (NO DECIMALS!)

Z = atomic # (# of protons & electrons)

X = symbol of the element

Isotopes

• Isotopes– Atoms with the same number of protons but a

different number of neutrons

– Atoms of the same element have the same atomic number but different mass numbers

Isotopes cont.

Potassium-39 Potassium-40 Potassium-41

Protons 19 19 19

Neutrons 20 21 22

Electrons 19 19 19

Isotopes cont.

Isotopes cont.

The diagram shows three oxygen isotopes. Each nucleus has eight protons (gray) and eight, nine, or ten neutrons (green).

Oxidation Numbers

• Indicate the charge on the ion

• Found on the periodic table

• Common oxidation numbers are given in additional table

• Note:– Group 8 elements do not form ions; No ox #– Transition metals have multiple ox #’s

Valence Electrons

• Electrons in the outer most shell

• The electrons on an atom that can be gained or lost in a chemical reaction

• More on this to come with Electron Configurations….

Alkali Metals

• Group IA (except H)• Li, Na, K, Rb, Cs, Fr• Soft, gray metals• Very reactive

– Especially with water

• React with water to form bases• “Alkali” = basic

Alkali Metals cont.

• Why are they so reactive?– One electron in their outer shell

• Only 1 electron away from a full outer shell

• Want to lose that electron: easily react

– So reactive – don’t occur as free elements

Interesting Tidbits

• Li – used as depression medication

• Cs – used in atomic clocks

• Fr – predicted by Mendeleev in 1870s; discovered in 1939

– Less than 1 oz. of Fr exists at any given time

Alkali Earth Metals• Group IIA• Be, Mg, Ca, Sr, Ba, Ra• Shiny, silvery-white metals• Harder and denser than Group IA• Distributed in rock formations• Reactive but not as reactive as Group IA

– 2 outer electrons

– Want to lose 2 to have a complete outer shell• +2 oxidation number

Interesting TidbitsInteresting Tidbits

• Used in pyrotechnics and fireworksUsed in pyrotechnics and fireworksMg – white; Sr – red; Ba – greenMg – white; Sr – red; Ba – green

Calcium

• Widely distributed as limestone• Important biologically for bones and teeth

• Compounds of Calcium- CaCO3 = limestone]- CaO = “lime” or “quicklime”- Ca(OH)2 = “limewater”; treat antacid

Nobel Gases

• Group VIIIA• He, Ne, Ar, Kr, Xe, & Rn• Aka “Inert Gases” b/c they are unreactive• Aka “Rare Gases” b/c they are very rare on Earth•Colorless, tasteless, odorless

Nobel Gases cont.• Why are they unreactive?

– Their outer shells are full– Recall: the outer shell electrons are the ones

involved in bonding– When the outer shells are full, these electrons can’t

bond and, therefore, react with other elements• Used in:

– Lighting • Fill light bulbs, neon lights, black lights,

flashlight bulbs, strobe lights, headlights, etc.

Halogens

• Group VIIA• F, Cl, Br, I, At• Non-metals

- Exist in all 3 states at room temperature:

* Solid: I, At* Liquid: Br* Gas: F, Cl

• Very reactive• Most often, bond with metals

•Diatomic (F2, Cl2, etc.)

Halogens cont.

• Why are the halogens diatomic?– 1 electron away from a full outer shell– Too reactive/unstable by itself– Bonds with another atom so both have 8– Dot diagrams

• What is the mole weight of chlorine gas?• Which is more reactive, F2 or Cl2?

– F2

Transition Metals

• Middle section of the periodic table

• Exhibit metallic properties– Ductile– Malleable– Good conductors of heat and electricity– Silvery luster (except Cu and Au)

Ions

• Atoms that have a positive or negative charge• To become an ion, an atom gains or loses

electrons• Cation = positively charged ion;

– Lose electron(s) – Metals form positive ions

• Anion = negatively charged ion;– Gain electron(s)

Polyatomic Ions

• Def: tightly bound groups of atoms that behave as a unit and carry a charge

• Ion composed of more than one atom• List of common polyatomic ions• Ex:

– SO4-2

– NO3-1 ** Must memorize these

– CO3-2 polyatomic ions! **

– PO4-3

– OH-1

– C2H3O2-1

– NH4+1