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2. Atoms, molecules, and ions

Early history of chemistry

prior to 1000 BCE

natural ores processed to produce metals for ornaments and weapons

use of embalming fluids

etc.

Greeks

400 BCE: all matter composed of four fundamental substances: fire, earth, water, air

whether matter is continuous (infinitely divisible into smaller pieces) or composed of small, indivisible particles

latter position: Demokritos/Democritus of Abdera (c. 460-370 BCE), Leucippos; used the term “atomos” (later became atoms) to refer to the particles

no definitive conclusion because no experiments

alchemy (pseudoscience)

some were mystics and fakes who were obsessed with turning cheap metals into gold

many were serious scientists; discovered several elements and learned to prepare mineral acids

16th century

Georg Bauer: development of system metallurgy (extraction of metals from ores)

Paracelsus: medicinal application of metals

Robert Boyle: first “chemist” to perform truly quantitative experiments

measured relationship between pressure and volume of air

published The Skeptical Chemist (1661)

ideas about chemical elements

no preconceived notion about number of elements

substance was element unless it could be broken down into two or more simpler substances

idea became generally accepted; Greek system of four elements eventually died

not always right; clung to alchemists’ views that metals were not true elements and that a way could be found to change one into another

combustion in 17th and 18th centuries

Georg Stahl: suggested that “phlogiston” flowed out of a burning material; substance in closed container stopped burning because the air became saturated with phlogiston

Joseph Priestley: oxygen gas supported vigorous combustion and therefore thought to be low in phlogiston

oxygen originally called “dephlogisticated air”

oxygen first discovered by Karl W. Scheele but his results were published after so Priestley is commonly credited with its discovery

Fundamental chemical laws

by late 18th century:

combustion studied extensively

carbon dioxide, nitrogen, hydrogen, oxygen discovered

list of elements continued to grow

Antione Lavoisie: finally explained true nature of combustion

regarded measurement as essential operation of chemistry

carefully weighed reactants and products of various reactions

law of conservation of mass: mass is neither created nor destroyed in a chemical reaction

combustion involved oxygen (Lavoisier named), not phlogiston

life supported by process involving oxygen; similar to combustion

first modern chemistry textbook: Elementary Treatise on Chemistry (1789); unified picture of chemical knowledge assembled up to that time

French Revolution → associated with collecting taxes for government → executed by guillotine as enemy of people

Joseph Proust: constant composition of compounds

law of definite proportion: a given compound always contains exactly the same proportion of elements by mass

John Dalton: atoms are particles that compose elements

if elements were composed of tiny individual particles, a given compound should always contain the same combination of those atoms

explained why same relative masses were always found

law of multiple proportions: when two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 g of the first element can always be reduced to small whole numbers

Dalton’s atomic theory

each element is made up of tiny particles called atoms

the atoms of a given element are identical; the atoms of different elements are different in some fundamental way(s)

chemical compounds are formed when atoms of different elements combine with each other; given compound always has same relative numbers and types of atoms

chemical reactions involve reorganization of atoms (changes in the way they are bonded together); atoms themselves not changed

atomic weights

8 g of oxygen known to be present for every 1 g of hydrogen in water

formula for water assumed to be as simple os possible; OH (hydrogen had mass of 1, oxygen had mass of 8)

created first table of atomic masses/atomic weights: weighted average mass of atoms in naturally-occurring element

many of masses later proved wrong due to incorrect assumptions

construction was important step forward

keys to determining formulas for compounds

Joseph Gay-Lussac: combining volumes of gases

2 volumes hydrogen + 1 volume oxygen → 2 volumes gaseous water

1 volume hydrogen + 1 volume chlorine → 2 volumes hydrogen chloride

Amadeo Avogadro

Avogadro’s hypothesis: at the same temperature and pressure, equal volumes of different gases contain the same number of particles

volume of gas determined by number of molecules, not size of particles

2 molecules of hydrogen + 1 molecule of oxygen → 2 molecules of water

formula for water is H₂O, not OH

interpretations were not accepted by most chemists; half-century of confusion

measurements of masses of elements that combined to form compounds

list of relative atomic masses

Jöns Jakob Berzelius: discovered cerium, selenium, silicon, thorium; developed modern symbols for elements used in writing formulas

diatomic molecules: H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂

Early characterization of atom

electrons

J. J. Thomson: electrical discharges in cathode-ray tubes (partially-evacuated tubes)

when high voltage applied, cathode ray (emanated from cathode, or negative electrode) produced; repelled by negative pole of applied electric field

postulated that ray was stream of negatively charged particles (now electrons)

charge-to-mass ratio of electron

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e represents charge on electron in coulombs (C)

m represents electron mass in grams (g)

determined by measuring deflection of beam of electrons in magnetic field

hoped to gain understanding of structure of atom

electrons could be produced from electrodes of different metals → all atoms must contain electrons

atoms electrically neutral → atoms must contain some positive charge

plum pudding model: diffuse cloud of positive charge with negative electrons embedded randomly

Robert Millikan: experiments with charged oil drops

determined magnitude of electron charge

calculated mass of electron

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radioactivity

late 19th: certain elements produce high-energy radiation

Henri Becquerel

piece of mineral containing uranium could produce its image on photographic plate in absence of light

radioactivity: spontaneous emission of radiation

three types of radioactive emission

gamma (γ) rays: high-energy “light”

beta (β) particles: high-speed electron

alpha (α) particles: 2+ charge (twice of electron and opposite sign); mass is 7300x of electron

nuclear atom

Ernest Rutherford: directed α particles at thin sheet of metal foil; particles should travel with little deflection

most particles went straight through, but many were deflected at angles, and some were reflected

proved plum pudding incorrect

deflection only possible through concentrated positive charge that contains most of mass

most α particles passed through foil because atom is mostly open space

nuclear atom: atom with dense center of positive charge (nucleus) with electrons moving around nucleus at large distance relative to nuclear radius

Modern atomic structure introduction

chemistry of atom comes mainly from electrons, so a crude model is sufficient

nucleus

assumed to contain:

protons: positive charge equal in magnitude to atom’s negative charge

neutrons: virtually same mass as proton but not charge

small compared to overall size of atom; extremely high density (almost all of mass)

nucleus size of pea would have mass of 250 million tons

different atoms have different chemical properties due to number and arrangement of electrons

electrons are the parts that “intermingle” when atoms form molecules

atoms of different elements (different numbers of protons and electrons) show different chemical behavior

isotopes: atoms with the same number of protons but different numbers of neutrons

symbol:

atomic number (number of protons): subscript

mass number (total number of protons and neutrons): superscript

almost identical chemical properties; most elements in nature contain mixtures of isotopes

Mass and charge of subatomic particles

Mass and charge of subatomic particles

Particle

Mass (kg)

Charge

electron

9.109 x 10^-31

proton

1.673 x 10^-27

neutron

1.675 x 10^-27

none

There are no rows in this table

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Molecules and ions

chemical bonds: forces that hold atoms together in compounds

covalent bonds: type of bonding where electrons are shared by atoms

molecule: collection of atoms

representing molecules

chemical formula

symbols for elements indicate types of atoms present

subscripts indicate relative numbers of atoms

structural formula

individual bonds are shown (lines)

dashed: behind plane of paper

wedge: in front of plane of paper

may or may not indicate actual shape of molecule

space-filling model

relative sizes of atoms

relative orientation in molecule

ball-and-stick model

three-dimensional position of atoms and bonds

atoms: spheres

bonds: rods

ion: atom or group of atoms with net positive or negative charge

when electron is transferred, atoms left with charge

cation: positive ion (Na → Na⁺ + e⁻)

anion: negative ion (Cl + e⁻ → Cl⁻)

ionic bonding: force of attraction between oppositely-charged ions

ionic solid: solid consisting of oppositely-charged ions

polyatomic ion: ion containing multiple atoms

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Periodic table introduction

periodic table: chart that shows all known elements and gives information

metals (majority)

physical properties

efficient conductors of heat and electricity

malleable (can hammer into thin sheets)

ductile (can be pulled into wires)

(often) lustrous

chemically: tend to lose electrons to form positive ions

nonmetals (fewer; upper-right except H)

lack physical properties that characterize metals

chemical properties

tend to gain electrons to form negative ions

often bond to each other by forming covalent bonds

elements in the same vertical columns (groups/families) have similar chemical properties

alkali metals (Group 1A): very active; form ions with 1+ charge

alkaline earth metals (Group 2A): form ions with 2+ charge

halogens (Group 7A): form diatomic molecules; form ions with 1- charge (except At)

noble gases (Group 8A): exist under normal conditions as monatomic gases; little chemical reactivity

symbols: typically 1A through 8A; sometimes 1 through 18

periods: horizontal rows (first, second, etc.)

Naming simple compounds

binary compound: two-element compound

binary ionic compounds

cation named first, anion second

cation: name from element

Roman numeral to indicate charge of cation (mostly for transition metals): Type II

elements that form only one cation do not need to be identified by a Roman numeral: Type I

special cases (no Roman numeral)

silver: almost always found as Ag⁺; typically called silver chloride instead of silver(I) chloride

zinc: only forms Zn²⁺ ion

anion

monatomic: root of element name + -ide

polyatomic: special names that must be memorized

oxyanion: anion consisting of an atom of an element and a number of oxygen atoms

-ite: smaller number of oxygen atoms

-ate: larger number of oxygen atoms

hypo-: fewest oxygen atoms

per-: most oxygen atoms

binary covalent compounds (two nonmetals): Type III

first element named first with full element name

second element named like anion

prefixes to denote numbers of atoms

mono- never used for first element

acids: substances that produce a solution containing free H⁺ ions (protons) when dissolved in water

if name of anion ends in -ide: root of anion name with prefix hydro- and suffix -ic

if anion contains oxygen

if name ends in -ate: root of anion name with suffix -ic

if name ends in -ite: root of anion name with suffix -ous

Common polyatomic ions

Common polyatomic ions

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Prefixes to indicate number in chemical names

Prefix

Number

mono-

di-

tri-

tetra-

penta-

hexa-

hepta-

octa-

nona-

deca-