Chem Exam 1

• Scientific method

• Process of studying natural phenomena involving observations, forming laws and theories, and testing theories by experimentation

• (Observation → hypothesis → hypothesis) → (Theory → Prediction → Experimentation)

 Qualitative observations

• Descriptive of qualities, characteristics, etc..

• Color, smell, appearance

• Quantitative observations

• Deals with specific numbers and values

• Volume, speed, time, mass

• Hypothesis

A possible explanation for an observation

• Experiment

Carried out to test the hypothesis. New info gathered and process repeats itself

Theory (model)

• Set of assumptions to explains some aspect of observed behavior; explain why nature behaves in a particular way and often change

• Natural law

A concise statement expressing generally observed behavior; implemented when same observation applies to many different systems

• Law vs theory

• A law is a summary of what happens while a theory is an attempt to explain why it happens

Matter

• Material if the universe, anything occupying space and having mass

Solid

Rigid, fixed volume and shape, slightly compressible

Liquid

• Definite volume, no specific shape (shape of container), slightly compressible

• Gas

• No fixed volume or shape (volume and shape of container), highly compressible (can pack a lot of gas into a small container)

• Pure substances

• A substance with a constant composition

• Water, sugar

• Compound

• A substance that can be further broken down into its constituent elements by chemical processes

• H2O → H, O  LiCl → Li, Cl

• Elements

• A substance that cannot be further broken down

• Anything on the periodic table

Most elements in nature as individual atoms are…

• Monatomic; unless otherwise noticed assume elements are monatomic

Many elements exist in their most stable state

H2, N2, O2, Cl2, Br2, I2, P4, S8

• Molecule

2 or more atoms held together by chemical bonds

All compounds are…

• Molecules, but not all molecules are compounds

• Mixture

• A substance with a variable composition

• Wood, gasoline, soil

• Homogeneous mixture

• Having visible indistinguishable part

• Salt and water, can’t see salt after mixing

• Heterogenous mixture

• Having visibly distinguishable parts

• Sweet tea

• Distillation

• Method of separating mixture that depends on boiling points  of a liquid mixture to separate

• Good to use if components of mixture have different boiling points or if samples dissolve in one another

• Filtration

• Mixture is passed through as mesh allowing liquid to pass and solid to stay

• Good to use if solid is suspended in liquid/ at bottom of container

• Law of conservation of mass

• Mass is neither creator nor destroyed

• Law of definite proportions/constant composition

• A given compound always contains the same proportion of elements by mas

• Ex: 3 samples of H2O all have 2 H and one O atoms

• Law of multiple proportions

• When 2 elements form a series of compounds, the ratios of the masses of the 2nd element combine with a fixed mass of the 1st element can always be reduced to small whole numbers

• The element with fixed mass is considered the first element

• Extensive property

• A property that depends on the amount of matter in a sample

• Mass, volume, length

• Intensive property

• Property that is independent of the amount of matter in a sample

• Color, melting point, boiling, point, boiling point, density, molar mass

• Density

• The mass of a substance per unit of volume

• d= mass/volume

• Density is why why certain items will float or sink, ex: ice → less dense, water → more dense

• Physical change

• Change in the form of a substance, not in the chemical composition

• Ice melting, water boiling off, cutting a sheet of paper

• Chemical change

• Change in the chemical composition of a substance

• Iron rusting, wood burning, baking a cake

Physical property

Odor, taste, appearance, melting/boiling point, conductivity, density, etc

• Chemical property

Corrosiveness, flammability, reactivity

• Law of conservation of energy

Energy cannot be created or destroyed, but it can be transferred; the total energy of the universe is constant

Potential energy

• Energy due to position or composition

• Ex: ball sitting on top of a hill, compressed spring

• Kinetic energy

• Energy due to motion

• Ex: Ball rolling down a hill, a spring that has been released

• Prefixes

• Giga- → G → 10^9 (billion)

• Mega- → M → 10 ^6 (million)

• Kilo- →  K → 10^3 (thousand)

• Deci- → d → 10^-1 (tenth)

• Centi- → c → 10^-2 (one hundredth)

• Milli- → m → 10^-3 (one thousandth)

• Micro- → μ → 10^-6 (one millionth)

• Nano → n → 10^-9 (one billionth)

• Pico → p → 10^-12

• Femto- → f → 10^-15

• If something doesn’t have a prefix

• 10^0

• To convert between metric prefixes…

• Take the exponential notation of what you are starting with and subtract it from what you are trying to get

• Ex: 4 GL to L → 4 GL x (10^9-0)= 4 x 10^9 GL

• 1cm^3 = 1 mL

• When taking a measure…

• Record all the certain digits plus the first uncertain digits

• Accuracy

• The agreement of a particular value with the true value

• Precision

• Refers to the degree of agreement among several measurements of the same quantity; looks at reproducibility; something can be precise but not accurate

• Sig figs

• Nonzero integers always counts as sig figs

• Preceding/leading zeroes are not significant Ex: 0.0053 → 2f

• Zeroes between nonzero digits are always significant

• Zeroes at the end of a number are only significant if the number contains a decimal point

• Exact numbers or things that can be counted have infinite sig figs as well as conversion factors (1 in = 2.54 cm), atomic and molar masses, and values in equations (temp. conversion)

• Measure values are not exact (12.00 meters, 25.00 degrees celsius)

For multiplication and addition…

We use the smallest number of sig figs

For addition and subtraction

We report our final value using the value with the smallest number of decimal places

If performing a series of calculations…

• Wait to round until the very end

Dimensional analysis/unit factor method

• To convert from one unit to another

• Boiling point of water

• 212 degrees F

• 100 degrees C

• 373.15 K

• All have infinite number of sig figs

• Freezing point of water

• 32 degrees F

• 0 degrees C

• 273.15 K

• All have infinite number of sig figs

• Nucleus

• Contains protons and neutrons; accounts for majority of the mass of the atom

• Proton mass is…

• Similar to neutron mass; both much higher than electron mass

• Number of protons is…

• Equal to number of electrons in neutral atoms

• Ions

• Form when electron number is changed

• Cation- a positively charged ion

• Anion- a negatively charged ion

• Polyatomic ion- an ion containing many atoms

• Nuclear symbol format

• Mass number/atomic number element symbol   or barium~138

• Mass number

• Protons and neutrons; not in the periodic table

• Atomic number

Just protons; in the periodic table

• Isotopes

• Atoms with the same number of protons and different neutrons

• Average atomic mass

• Consider relative abundance of each isotope of an element

• C→ (12.00 x 98.89%) + (13.0034 x 1.11%) = 12.01 u

• If one single atom was isolated it would never weigh 12.01 since it is an average

• Dalton’s postulate/ Dalton’s atomic theory part 1

• Each element is made up of atoms. Atoms are indivisible → no longer accurate since atoms are made of p+, n, e-, can be split

• Dalton’s postulate/ Dalton’s atomic theory part 2

Atoms of a given element are identical; the atoms of different elements are different in some fundamental way → discovery of isotopes proved that not all atoms of an element are identical

• Dalton’s postulate/ Dalton’s atomic theory part 3

• Chemical compounds form when different elements combine with each other. A given compound always has the same relative number and types of atoms → stated by law of constant composition

• Dalton’s postulate/ Dalton’s atomic theory part 4

• Chemical reactions involve reorganization of atoms - changes in the way they are bound together. Atoms themselves are not changed in a chemical reaction → sometimes they do change 

• Thomson’s Cathode Ray Experiment

• Thompson used a cathode ray tube to create negatively charged particles (electrons)

• Electrons were repelled by the negative electric field, indicating they were negatively charged

• He believed atoms were neutrally charged, so if a negative charge was present there had to be a positive charge

• Led to “plum pudding model”

• Rutherford Gold-Frail experiment

• Used to test out the “plum pudding” model

• Focused beam of a particles (positively charged) at gold foil; proved “plum pudding” model was incorrect

• Experiment confirmed presence of nucleus since some a particles scattered

• Metals

• Tend to lose electrons to form positive ions

• Conducts electricity well

• Malleable and ductile

Nonmetals

• Tend to gain electron to form negative ions

• Variety of properties; missing properties found in metals