CHEM EXAM 1: Concepts, Definitions, and other facts

Matter

anything that has mass and occupies space

Fact

A basic statement established by experiment or observation that is true under specific conditions of the observation

Hypothesis

tentative(can be changed) explanation that can be tested by further investigation

Theory

Well-supported explanation of observations, aren’t likely to change

Law

Principles that can be used to predict the behavior of the natural world (patterns)

Qualitive vs Quantitative Observations

• Qualitive observations describe using senses

• Quantitative observations are measured, include number values

Scientific Method

1. Make observations

2. Formulate a hypothesis

3. Test the hypothesis through experimentation

4. Accept or modify the hypothesis

5. Develop into a law and/or a theory

Scientific Notation

N ×10ⁿ where N is less than 10 and n is any integer

Sig Fig rule when multiplying/dividing

count the number of sig figs in each factor and limit the sig figs in the answer to the least # of sig fig in the factors (use rounding)

Sig Fig rule when adding/subtracting

limit the sig figs in the answer to the digits that both numbers have in common

accuracy vs precision

• accuracy is how close a measurement is to the true value

• precision is how close a series of measurements are to each other and if hey give reproducible results

Sig Fig Rules

1. All nonzero digits in a measurement are significant

2. Zeroes that appear between other nonzero digits (middle zeros) are always significant

3. Zeros that appears in front of all the nonzero digits (leading zeros) are never significant

4. Zeros that appear after all nonzero digits (trailing zeros) are ambiguous

5. Zeros after the decimal pt are significant, but not if they are leading zeros

Mass vs Weight

• Mass measures amount of matter

• Weight is the force of gravity

Volume

amount of space occupied by matter

Temperature

measure of the average amount of energy of motion (kinetic energy) a system contains

Density

the mass of an object divided by its volume; can be used as a conversion factor between mass and volume

Atoms

building blocks of matter and the universe

Molecules

bonded atoms, can be combined in different ways, have properties like shape, size, color, boiling point, volatility, conductivity, etc.

Stages of Matter

• Solid: particles are in a fixed position, shape and volume are definite

• Liquid: particles are in contact but not fixed, shape is not definite by volume is

• Gas: particles are in random positions, shape and volume aren’t definite

Pure substance

matter that has constant composition and properties are constant throughout the sample

Mixture

matter consisting of two or more substances that retain their individual identities and can be separated by physical methods

• Homogenous mixture: mixtures that have a uniform composition and properties throughout (ie AJ and tea)

• Heterogeneous mixture: mixtures that aren’t uniform (ie guac and trail mix)

Element

substance that cannot be broken down into chemically simpler components; composed of the same type of atom

Compound

substance that are made up of more than one type of atom; can be separated into simpler substances and elements by chemical methods

Physical vs Chemical Properties

• Physical properties: characteristics that can be observed or measured

  • Ex: conductivity, malleability, color, hardness, solubility, density

• Chemical properties: describe a sample’s potential to undergo a chemical reaction by virtue of its composition

  • Ex: burn, rot, explode, decompose, ferment, rust, flammability

Physical vs Chemical Change

• Physical: no bonds are broken or formed

  • Ex: change of state, separation of a mixture, deformation, making solutions

• Chemical: bonds are broken and/or formed

  • Indicators: temp changes, light produced, color changes, bubbles, different smell or taste, precipitate

Conservation of mass

No matter can be created or destroyed

Heat

energy that is transferred from a hot object to a cooler object due to the difference in their temperature

Work

transferred energy as a result of a force applied over distance

Kinetic and Potential Energy

• Kinetic: energy of motion

• Potential: stored energy that depends on the position of an object rea=lative to another object

Endothermic Processes

Where a system absorbs heat and the temperature of the surroundings get cooler; the system gets hotter (q>0)

Exothermic Processes

Where a system releases/loses heat and the temperature of the surroundings get warmer; the system gets colder (q<0)

Heat Capacity vs Specific Heat

• Heat capacity: the amount of heat required to raise the temperature of an object by 1 °C

• Specific Heat: the amount of energy required to raise the temperature of 1 gram of the substance by 1 °C

Diatomic Atoms

• Hydrogen (H)

• Oxygen (O)

• Fluorine (F)

• Bromine (Br)

• Iodine (I)

• Nitrogen (N)

• Chlorine (Cl)

John Dalton and Modern Atomic Theory

• all matter is composed of extremely small atoms

• atoms of a given element are identical in size, mass, and other properties

• atoms cannot be subdivided, created, or destroyed

• atoms of different elements can combine in numerous rations to form chemical compounds

• in chemical reaction, atoms are combined, separated, or rearranged

Plum Pudding model by J.J. Thompson

electrons are like blueberries on a muffin or plums in pudding embedded in a uniform sphere of positive charge

Nuclear Model by Rutherford

• all of the positive charge and majority of the mass of the atom must be concentrated in the atom’s nucleus

• Nucleus is the central core of the atoms that is composed of protons and neutrons

• Electrons are distributed around the nucleus and occupy most of the volume of the atom

Electrons

• contribute virtually nothing to the total mass of an atom

• found orbiting nucleus

• charge of -1

• symbol: e^-

Protons

• found in the nucleus

• charge of +1

• mass of 1 amu

• symbol: p⁺

Neutrons

• in all atoms except hydrogen

• no charge

• mass of 1 amu

• symbol: n⁰

Periodic Table

• Developed by Dmitri Mendeleev

• Rows of the table are called periods

• Columns of the table are called groups

• Each element has unique symbol and atomic number

Metal

a substance that is shiny, an excellent conductor of electricity and heat, and malleable and ductile

Metalloid

elements with properties intermediate between those of metals and nonmetals

Nonmetal

dull, poor conductor, and brittle

Group 1

The Alkali Metals; common in nature and daily life

• Lithium (Li), atomic # of 3

• Sodium (Na), atomic # of 11

• Potassium (K), atomic # of 19

• Rubidium (Rb), atomic # of 37

• Cesium (Cs), atomic # of 55

• Francium (Fr), atomic # of 87

“Lisa Saw Patrick Rob Cane’s Fries”

Group 2

The Alkaline Earth Metals

• Beryllium (Be), atom # of 4

• Magnesium (Mg), atom # of 12

• Calcium (Ca), atom # of 20

• Strontium (Sr), atom # of 38

• Barium (Ba), atom # of 56

• Radium (Ra), atom # of 88

“Bart Might Cause Some Bad Rukus”

Group 7

Halogens; react readily with metals to form compounds

• Florine (F), atomic # of 9

• Chlorine (Cl), atomic # of 17

• Bromine (Br), atomic # of 35

• Iodine (I), atomic # of 53

• Astatine (At), atomic # of 85

“Flanders Can Bake Incredible Apple-pie”

Group 8

Noble Gases; compromised of single atoms aka monatomic; unreactive

• Helium (He), atomic # of 2

• Neon (Ne), atomic # of 10

• Argon (Ar), atomic # of 18

• Krypton (Kr), atomic # of 36

• Xenon (Xe), atomic # of 54

• Radon (Rn), atomic # of 86

“Home Needs All Krustry’s Extras Ribs”

Ions

Atoms that have a charge because they have more protons than electrons pr vice versa

• cations: have a positive charge because they lost an electron; metals become cations (groups 1 & 2 become cations)

• anions: have a negative charge because they gained an electron; nonmetals become anions (groups 7 become anion)

Isotopes

• Atoms of the same element that differ in their number of neutrons

  • have different mass numbers

• can be denoted like so:

  • Nickel-59 where # is the mass number

Mass Number (A)

total number of protons and neutrons

Chemical Potential Energy

Potential Energy stored in atoms

Physical Chemistry

Study of macroscopic properties, atomic properties, and phenomena in chemical systems

Organic Chemistry

Study of chemicals containing carbon

Inorganic Chemistry

Study of chemicals found in rocks and minerals

Analytical Chemistry

Study of composition of matter through separation, identification, and quantification of chemicals in samples of matter

Biochemistry

Study of chemical processes that occur in living systems

Alchemy

Study of matter based on fire, water, earth, and air that believed changing proportions of these in a substance could change its composition

Robert Boyle

Developed basic ideas about behavior of gases

Joseph Priestly

Isolated and characterized Oxygen, Carbon Dioxide, etc.

C. W. Scheele

Discovered Chlorine gas

Lavoiser

• Father of chemistry

• discovered nitrogen gas

• discovered the role of oxygen in combustion

• formulated the law of conservation of matter

Amadeo Avogadro

Pioneered quantitative approach to chemistry by calculating the number of particles in a given amount of gas