1St Semester Final Exam Review Material

1st semester Final Exam review material.

Unit 1: Measurement and Calculations

What is Chemistry?

Matter is anything that has mass and occupies space.
You don’t have to be able to see something for it to qualify as matter
Chemistry is the study of the composition of matter and the changes that matter undergoes.
Because living and nonliving things are mad of matter, chemistry affects all aspects of life and most natural events.

The scientific Method

Scientific method is a logical, systematic approach to the solution of a scientific problem
Steps in the scientific method include making observations, testing hypotheses, and developing theories.

1. Making Observations

When you use your senses to obtain information you make an observation.

Ask a question The observation can lead to a question
Example: observation = flashlight won’t come on, question = what’s wrong with the flash light?

3. Developing a hypothesis

A hypothesis is a proposed explanation for an observation.
Example: guessing the batteries are dead would be a hypothesis

4. Testing Hypotheses

Experiment is a procedure that is used to test a hypothesis
independent variable is the variable that you change during an experiment
dependant variable is the variable that is observed during the experiment.
For the results of an experiment to be accepted they MUST be able to be reproduced.

5. Analyze Data

Look at the information from your test to see if your hypothesis is correct
If correct: reinforce hypothesis with additional experimentation
If incorrect: reevaluate the hypothesis and create a new experiment

6. Developing Theories

a hypothesis becomes a theory when it meets the test of repeated experimentation.
Theory is a well-tested explanation for a broad set of observations.
Theories can never be proven; they may be changed at some point in the future to explain new observations or experimental results.

. Scientific Laws

Scientific law is a concise statement that summarized the results of many observations and experiments
A scientific law does not attempt to explain why something occurs

Using and Expressing Measurements

Scientific notation, a short way of writing large or small numbers, a given number is written as the product of two numbers: a coefficient and 10 raised to a power.
Example: 602,000,000,000,000,000,000,000 would be written as 6.02 x 1023

Accuracy, Precision and Error

Accuracy is a measure of how close a measurement comes to the actual or true value of whatever is measured.
Precision is a measure of how close a series of measurements are to one another
To evaluate the accuracy of a measurement, the measured value must be compared to the correct value. To evaluate the precision of a measurement, you must compare the values of two or more repeated measurements.
The center is the true value, A has both accuracy (near center) and precision (darts close to one another. B has only precision because darts are close to another but not the center. C has no accuracy or precision.
Accepted values is the correct value based on reliable references
Experimental value is the value measured in the lab.
Error is the difference between the experimental value and the accepted value
Error can be positive or negative
Percent error is the absolute value of the error divided by the accepted value, multiplied by 100%

Significant Figures in Measurements

Significant figures include all of the digits that are known, plies a last digit that is estimated
Measurements must always be reported to the correct number of significant figures because calculated answers often depend on the number of significant figures in the values used in the calculations.

Rules for determining whether a digit in a measured value is significant

Nonzero digits are significant. 5.23 has 3 significant figures
Zeros between nonzero digits are significant. 5001 has 4 significant figures
Zeros in front of nonzero digits are not significant, they are only place holders. 0.000099 has 2 significant figures
Zeros at the end of a number and to the right of a decimal place are significant. 1.0100 has 5 significant figures
Zeros to the left of an understood decimal point are not significant, they are only place holders. 55000 has 2 significant figures
Defined quantities and counted quantities have unlimited number of significant figures.

Significant figures in Calculations

In general, a calculated answer cannot be more precise than the least precise measurement from which it was calculated.
When rounding first decide how many significant figures the answer should have.
Next round to that number of digits, counting from the left.
If the number to right of the last significant digit is 4 or less round down, if it is 5 or up round up.
Example: each rounded to 3 significant figures 5.236 = 5.24, 8.023 = 8.02
With addition or subtraction the calculation should be rounded to the same number of decimal places (NOT digits) as the measurement with the lease number of decimal places
Example: [2.01 has the lease number of decimal places]
With multiplication and division the calculation should be rounded to the same number of significant figures as the measurement with the lease number of significant figures.
Example: [12 has only 2 significant figures]

Units and Quantities

Meter is the SI basic unit of length
Common metric units of length include the centimeter, meter, and kilometer
Scientists commonly use two equivalent units of temperature, the degrees Celsius and the Kelvin.

Conversion Factors

Many quantities can usually be expressed different several different ways.
Example: 1 dollar = 4 quarters = 10 dimes = 100 pennies
Whenever two measurements are equivalent, a ratio of the their measurement will equal 1
Conversion factor is a ratio of equivalent measurements.
Example:
When a measurement is multiplied by a conversion factor, the numerical value is generally changed, but the actual size of the quantity measure remains the same.
Example: 2 hours = 120 minuets = 7200 seconds

Metric Units and Values

Prefix / tera / giga / mega / kilo / hecta / deca / base / deci / centi / milli / micro / nano / pico
symbol / T / G / M / k / h / da / d / c / m / μ / n / p
meaning / 1012 / 109 / 106 / 103 / 102 / 101 / 100 / 10-1 / 10-2 / 10-3 / 10-6 / 10-9` / 10-12
pneumonic / the / great / mad / king / Henry / died / by / drinking / chocolate / milk / under / nick's / porch

Metric conversion use the following basic equation:

Prefix units have 2 or more letters: kg, pm, TJ

Base units have 1 letter: m, g, L, J, N, s

Dimensional Analysis

Dimensional Analysis is a way to analyze and solve problems using the units of the measurements.
Dimensional analysis provides you with an alternative approach to problem solving.
Example: 9875 seconds equals how many hours?
The key to dimensional analysis is to set it up so that the UNITS cancel.

Determining Density

Density is the ratio of the mass of an object to its volume

Unit 2: Matter and change

Identifying Substances

Substance is matter that has a uniform and definite composition.
Gold and copper are examples of substances and are also referred to as pure substances
Every sample of a given substance has identical intensive properties because every sample has the same composition
Physical property is a quality or condition of a substance that can be observed or measured without changing the substance’s composition.
Physical properties include hardness, color, conductivity, malleability, melting point and boiling point.
Physical properties can help chemists identify substances

2.1 C. States of Matter

1. Solid

Solid is a form of matter that has DEFINITE shape and volume
The shape of a solid doesn’t depend on the shape of a container
The particles in a solid are packed tightly together and often in an orderly arrangement
Solids have vibrational kinetic energy (the atoms vibrate around a fixed position)
Solid cannot be compressed

2. Liquid

Liquid is a form of matter that has a DEFINIE volume but NOT a definite shape
a liquid takes the shape of it’s container
The volume of the liquid doesn’t change as it’s shape changes, it is constant
the particle in a liquid are in close contact with one another but not as rigid or orderly as a solid
in a liquid the particle are free to flow past one another and have more space between the atoms
Liquids have vibrational and rotational kinetic energy
Liquids are slightly compressible

3. Gas

Gas is a form of matter that takes both the shape and volume of its container
The particles in a gas are much farther apart than the particles in a liquid.
Gases are easily compressed into smaller volumes
Vapor describes the gaseous state of a substance that is generally a liquid or solid at room temperature.
Gases have vibrational, rotational and translational kinetic energy
The atoms in gases have the most space between them

Physical Changes

Physical change a change during which some properties of a material change, but the composition of the material does not change
Boil, melt, freeze, condense, breaking, splitting , grind, and cut are used to describe physical changes.
Physical changes can be classified as reversible or irreversible.
Melting, freezing and boiling are examples of reversible physical changes (can be changed back)
Cutting, grinding, and breaking are example of irreversible physical changes (cannot be changed back)

A. Classifying Mixtures

Mixture is a physical blend of two or more components (substances)
Most samples of matter are mixtures
Based on the distribution of their components, mixtures can be classified as heterogeneous mixtures or as homogeneous mixtures.

1. Heterogeneous Mixture

Heterogeneous mixture is a mixture in which the composition is NOT uniform throughout.
Examples: salad, pizza, beach sand

2. Homogeneous mixture

Homogeneous mixture is a mixture in which the composition IS uniform throughout.
Example: vinegar, soda, tap water
Solution is a homogenous mixture where solutes are dissolved in a solvent (kool-aid)
most solutions are liquids but some can be solid (steel, bronze) and some are gases (air)
Phase describes any part of a sample with uniform composition and properties (solid, liquid, gas)
all homogeneous mixtures consist of a single phase

Distinguishing Elements and Compounds

A pure substance can be either an element or a compund
Element is the simplest form of matter that has a unique set of properties
Element cannot be broken down into simpler substances. Examples: oxygen, nitrogen, sodium
Compound is a substance that contains two or more elements chemically combined in a fixed proportion.
Compounds can be chemically separated in to simpler substances. Examples: water, sugar, oil
Compounds can be broken down into simpler substances by chemical means, but elements cannot.

1. Breaking Down Compounds

Chemical change is a change that produces matter with a different composition than the original matter
Heating a substance can be used to break down compound into simpler substances.
Example: heating sugar with give you carbon and water

2. Properties of Compounds

The properties of compounds are usually quite different from those of their component elements.
Carbon: black tasteless solid, sugar: white sweet solid
Hydrogen: flammable gas, Oxygen: color gas that supports burning, water: liquid that can stop materials from burning

Symbols and Formulas

Chemists use chemical symbols to represent elements, and chemical formulas to represent compounds
Chemical symbol is the one or two letter abbreviation for an element
The first letter of a chemical symbols is ALWAYS capitalized and the second letter is always lowercase.
Chemical symbols provide a short hand way to write the chemical formulas of compounds
Subscripts are used to show the number of atoms of a given element in a compound’s chemical formulas.
Example: C12H22O11 is the formula for table sugar, it has 12 Carbon atom, 22 Hydrogen atoms, and 11 Oxygen atoms.

Chemical Changes

Chemical property is the ability of a substance to undergo a specific chemical change
Words that signify a chemical change: burn, rot, rust, decompose, ferment, explode, and corrode
Chemical properties can be used to identify a substance
During a chemical change, the composition of matter always change

Recognizing chemical changes

Possible clues to chemical changes include a transfer of energy, a change in color, the production of a gas, or the formation of a precipitate
Precipitate is a solid that forms and settles out of a liquid mixture

Conservation of Mass

During any chemical reaction, the mass of the products is always equal to the mass of the reactants.
Mass is also held constant during a physical change
The law of conservation of mass states that in ANY physical change or chemical change the mass is conserved (stays the same).
Mass is neither created nor destroyed

Phase Change and Temperature

During a phase change the temperate of a substance remains the same. C= temperature that the substance is melting, G = temperature that the substance is boiling.

Chapter 4: Atomic Structure

Subatomic Particles

Three kinds of subatomic particles are electrons, protons, and neutrons.

1. Electrons

1897 J.J. Thomson discover electrons using the cathode ray tube
Electrons are negatively charged subatomic particles

2. Protons and Neutrons

Atoms have not net electric charge, they are electrically neutral
Electric charges are carried by particles of mater
Electric charges always exist in whole-number multiples of a single basic units
When a given number of negatively charged particles combines with an equal number of positively charged particles
Protons are positively charged subatomic particles
Neutrons are subatomic particles with not charge but with a mass nearly equal to that of a proton

The Atomic Nucleus

When subatomic particles were discovered, scientists wondered how these particle were put together in an atom
J.J. Thompson’s model was known as the ‘plum-pudding model”

1. Rutherford’s gold-foil Experiment

1911 Rutherford decided to test the current theory of atomic structure
He shot alpha particles (positive particle) at thin sheet of gold
The majority of the particles passed straight through
A small fraction bounced off the gold foil at very large angels

2. The Rutherford Atomic Model

Based on the experimental results Rutherford suggested a new theory of the atom
He suggested that the atom is mostly empty space (explaining the lack of deflection of most alpha particles)
Also that all the positive charge and mass is concentrated in a small region
Nucleus is the tiny central core of an atom and is composed of protons and neutrons.
In the nuclear atom, the protons and neutrons are located in the nucleus. The electrons are distributed around the nucleus and occupy almost all the volume of the atom

Atomic Number

Elements are different because they contain different numbers of protons
Atomic number of an element is the number of protons in the nucleus of an atoms in their element
Remember that atoms are electrically neutral
Because of that the number of electrons must equal the number of protons

Mass Number

Mass number is the total number of protons and neutrons in an atom.
If you know the mass number and the atomic number of any atom you can determine the number of neutrons in the atom
The number of neutrons in an atoms is the difference between the mass number and the atomic number
Short hand notation: , or gold-197

Isotopes

Isotopes are atoms that have the same number of protons but different number of neutrons
Because isotopes of an element have different numbers of neutrons, they also have different mass numbers.
Isotopes are chemically alike because they have identical numbers of protons and electrons (which are responsible for chemical behavior)

Unit 4: Modern atomic theory and Periodic Trends

Atomic Orbitals

An atomic orbital is often thought of as a region of space in which there is a high probability of finding an electron
Each energy sublevel corresponds to an orbital of a different shape, which describes where the electron is likely to be found.
s orbitals are spherical, the probability of finding an electron at a given distance from the nucleus in an s orbital does not depend on direction
p orbitals are dumbbell-shaped. The three kinds of p orbital’s have different orientations in space.
Four of the five kinds of d orbitals have clover leaf shapes. The shapes of f orbitals are more complicated than for d orbitals.
The numbers and kinds of atomic orbitals depend on the energy sublevel.
The lowest principal energy level (n = 1) has only one sublevel, called 1s.
The second principal energy level (n = 2) has two sublevels, 2s and 2p. the second principal energy level has four orbitals: 2s, 2px, 2py, and 2pz.
The third principal energy level (n = 3) has three sublevels. 3s, 3p, and 3d. Thus the third principal energy level has nine orbitals (one 3s, three 3p, and five 3d orbitals).
The fourth principal energy level (n = 4) has four sublevels, 4s, 4p, 4d, and 4f. The fourth principal energy level, then, has 16 orbitals (one 4s, three 4p, five 4d, and seven 4f orbitals).

Electron Configurations

In an atom, electrons and the nucleus interact to make the most stable arrangement possible.
The ways in which electrons are arranged in various orbitals around the nuclei of atoms are called electron configurations.
Three rules—the aufbau principle, the Pauli exclusion principle, and Hund’s rule—tell you how to find the electron configurations of atoms

1. Aufbau Principle

aufbau principle, electrons occupy the orbitals of lowest energy first.
The orbitals for any sublevel of a principal energy level are always of equal energy.
within a principal energy level the s sublevel is always the lowest-energy sublevel.
range of energy levels within a principal energy level can overlap the energy levels of another principal level.

2. Pauli Exclusion Principle

Pauli exclusion principle, an atomic orbital may describe at most two electrons.
To occupy the same orbital, two electrons must have opposite spins
A vertical arrow indicates an electron and its direction of spin (↑ or ↓).
An orbital containing paired electrons is written as

3. Hund’s Rule