Chem Chapters 1 & 2

Definition of Chemistry

the science or study of matter and changes in matter

Matter

has mass & occupies space (ex. dust, air, juice, molecules)

Not matter - energy forms & abstract things (ex. heat, light, & emotions)

Scientific Method

approach to gain knowledge

Scientific Law

collection of observations/measurements may show a general pattern of rule —> predict outcomes but don’t explain how or why

Hypothesis

tentative explanation or model to explain how and why a phenomenon occurs

educated guess based on known principles or logic then tested by experiment

Theory

well-established model or explanation that is testable & builds on body of knowledge

Macroscopic Domain

properties & observations made in the “big world”

ex. temperature, color,

Microscopic Domain

behavior explained by using atoms, ions, & molecules

ex. seen under microscope

Symbolic domain

use of formulas, equations, or graphs to describe properties & reactions

ex. atomic number, formula

Solid

fixed shape (does not flow), fixed volume, difficult to compress (ex. wood, clothing)

Liquid

fixed volume, difficult to compress, can flow, takes shape of bottom of container (ex. seawater, paint)

Gas

no fixed shape or volume, can flow, expands to fill container, easy to compress (ex. helium, air)

Plasma

like gas but electrically conducting, usually emitting light (ex. lightning, neon lamps)

Kinetic Molecular Theory (KMT)

chemical units (ions, atoms, molecules) in permanent motion based on temperature

Low Temperature (Low KE)

get solid —> attraction between particles dominate, motion limited to vibration in place, cannot slide past each other

Moderate Temperature

get liquid —> motion sufficient to open gaps, can flow by sliding past each other but attraction remains

High Temperature

get gas —> sufficient motion to overcome attractions, particles are free with large spaces, collisions

Extreme Temperature

get plasma —> enough energy to remove electrons from particles, leaving positive ions additional strong ectrostatic attractions & repulsions, emit light

Mass

amount of matter, inertial resistnace to force m=F/a (kg,g,lb, oz)

more matter = larger mass, harder to move & stop

independent of gravity & location

Weight

gravitational force on mass (N(newton) or lb-force)

depends on location

earth weight in mass units is the same as mass g=9.78 m/s²

Pure Substances

have fixed compositions, fixed properties

cannot be separated by physical means (ex. filtration)

elements or compounds

Elements

pure substances with only one kind of atom

Cannot be PHYSICALLY OR CHEMICALLY seperated (element = simplest possible)

Compounds

have two or more kinds of atoms BONDED in fixed ratios

CAN be chemically separated but CANNOT be physically separated

have fixed properties (unlike mixtures)

because of chemical bonding compounds are distinct form elements involved

Mixture

physical blend of 2 or more pure substances

retains some properties of substances involved

CAN be separated PHYSICALLY into simpler mixtures or substances

Heterogenous Mixtures

properties vary from place to place (“chunky or bubbly” materials)

most materials are this kind

Ex. sand or soil mixed with water, salad dressing, trail mix

Homogenous Mixture

properties are uniform throughout the mixture = solution

often called solutions especially for liquid solvents

appear uniform like pure substances but physically separable

ex. red wine, salt water, steel, brass

Colloids

look uniform but are heterogenous under a microscope (particles or bubbles are present)

ex. milk, blood, lotion, fog/cloud

How to Classify Matter

not uniform = heterogenous mixture

uniform under a microscope = homogenous mixture

1 substance = not mixture = pure substance

• only 1 atom = element

• else = compound

Atoms

smallest unit of ordinary matter and of an element

tiny nucleus of protons & neutrons surrounded by moving electrons

compounds involve whole number ratios of atoms of 2 or more

ex. He, O₂, Na

Molecules

discrete unit having 2 or more atoms bonded COVALENTLY together

1 kind of atom —> elemental molecules especially nonmetals, ex. N₂)

2 or more —> molecular compounds (majority of molecules) ex. CH₄

Law of Conservation of Mass/Matter

total mass after reaction = total mass before reaction in a closed system

matter/mass cannot be destroyed nor created in reactions (ex. old battery mass = new battery mass)

Law of Definite Proportions/Constant Composition

compound is always made of same element set in the same mass ratio regardless of source or synthesis method

Properties

characteristics of substance that help distinguid it from other substances

Physical Properties

can be observed without a change in composition (w/ out reaction)

observed/measured without change & same with physical change (same substance in different forms)

ex. mass, melting point, color, density, volume, hardness

Chemical Properties

those based on chemical reactions —> change in chemical identity

reactivity or inertness

ex. burning, acidity, toxicity, reduction potential

Extensive Properties

increase with the amount or extent of material

grows with sample

ex. mass, volume, surface area, moles

Intensive Properties

independent of size or extent of mass

ratios of extensive properties

color/density = mass/volume, specific heat = heat capacity & flammability

remains the same even if sample splits

Metals

good conductors of heat & electricity

shiny (lustrous), reflect light & heat

are plastic (can be shaped by force but not brittle)

• Malleability = flatten into a sheet

• Ductility = draw into wires

Non-metals

poor conductors (insulators) of heat & electricity

any state often gases, may be liquids or solids (solids are brittle)

top right of periodic table

Metalloids (semi-metals)

between non-metals & metals (on stair like bottom)

conductivity in between metals & nonmetals = semi conductive

all solids sometimes lustrous crystals

Measurement

has number, has units, has uncertainty

Chemistry is quantitative involves measurements of properties

SI Units

international system of units based on metric system

Length (SI Units)

meter, m

Mass (SI Unit)

kilogram, kg or gram, g(1000 g = 1 kg)

Time (SI Unit)

seconds, s

Amount of Matter (SI Unit)

mole, mol

Temperature (SI Unit)

kelvin, k

Electrical Current (SI Unit)

ampere, a

Luminous Intensity (SI Unit)

candela, cd

Density

intensive physical property, defined as mass/volume, represents how closely packed matter (ions, atoms, molecules)

d=m/v

floating & sinking dependent on density

Uncertain Digit Location

All measurements (except counting) involve uncertainty or error

use significant figures + unit measurements

• Sig Fig —> all certain digits & 1 uncertain digit

  • Usualy asumme ± 1 in last digit’s uncertainty

Sig Fig Counting

All nonzero digits significant

Zeroes between non zeroes significant —> 4.07 = 3 sig figs

Zeroes on let NEVER SIGNIFICANT (skip & count digits)

Zeros on right w/ decimal points significant

Rounding Rules

Look at discarded digits

IF discard <5, chop IF discard >5 chop and 1 to retained digit

Accuracy

closeness of result (or mean of set) to true/accepted value

Precision

closeness of repeated measurements to each other - how closely they agree

Early Ideas of Atomic Theory

Concept of atoms was fire proposed by Greek philosophers Leucippus & Democritus in 5th century BC

Aristotle believed that matter consisted of various combinations of the 4 elements - “fire, Earth, air, & water”

Dalton’s Atomic Theory - 5 Postulates

English teacher John Dalton proposed his atomic theory

1) Matter is composed of exceedingly small particles called atoms

2) An element consists of only one type of atom, which has a mass that is characteristic of the element & is the same for all atoms of that element

3) Atoms of an element differ in properties from atoms of all other elements

4) A compound consists of atoms in 2 or more elements combined in a small, whole number ration (always present in same ratio)

5) Atoms are neither created nor destroyed during a chemical change but instead rearranged to yield different types of matter

Law of Conservation of Matter

Lavoisier’s theory atoms are neither created nor destroyed during a chemical change, then the total mass of matter present when matter changes from one type to another will remain constant

Law of Definite Propionate or Law of Constant Composition

Prousts’s law that All samples of a pure compound contain the same elements in the same proportion of mass

Law of Multiple Proportions

Dalton’s law that when two elements react to form more than one compound, a fixed mass of one element will react with masses of the other element in a ratio of small, whole numbers

Electron Discovery

JJ Thompson experimented with cathode ray tubes that find electrons = negatively charged, subatomic particle with mass more than 1,000x less than that of an atom

Thompson able to calculate charge to mass ratio of cathode ray particles

Showed that the charge of mass raito of an electron 1.759×10¹¹ c/kg

Oil Drop Experiment

Robert A. Milikan’s created microscopic oil droplets, which were electrically charged

Concluded that 1.6 ×10^-19 was the charge of a single electron

Gold Foil Scattering Experiment

Ernest Rutherford’s expriment aimed a beam of alpha particles (a particles) at a very thin piece of gold foil

A small relatively heavy and + charged body, nucleus at the center of each atom

• Contains most of atoms mass

• Negatively charged electrons surround the nucleus

• Proton (+ charged) subatomic particles located in nucleus

Isotopes

atoms of the same element that differ in mass (Frederick Soddy)

Neutrons

uncharged, subatomic particles with mass approximately the same as that of protons

• Discovered by James Chadwick

• Also found in nucleus

Atomic Structure & Symbolism

The nucleus contains the majority of an atom’s mass

Protons & neutrons are much heavier than electrons

• Electrons occupy almost all of an atom’s volume

• Diameter of an atom = 10^-10

• Diameter of a nucleus = 10^-15

• Atoms & subatomic particles are very small

Proton mass = 1.0073 amu, charge +1

Neutron Mass = 1.0087 amu, charge 0

Electron Mass = 0.00055 amu, charge -1