Chemistry Exam 1 (Chapers 1-3.2) Study Guide

For Chem 161 Exam 1 (Except for polyatomic ions in a different set)

What is chemistry?

The study of matter and its formations

Matter

Anything that has mass and takes up space

Phases of Matter:

Solid, liquid, and gas

Solid Conditions

• Particles close together

• Constant volume

• Particles fixed in position

  • Definite shape

Liquid Conditions

• Particles close together

• Constant volume

  • Takes on shape of container

Gas conditions

• Particles far apart

• Will occupy any space available

Temperature

Measure of average kinetic energy

How do you change the state of matter?

Heat up or make it colder (take away heat)

Changes the vibration speed, causing phase change

Kinetic Energy=

(1/2)(mass)(velocity)²

Physical Properties

Description of matter

Qualitative or quantitative

It doesn’t change chemical composition

Physical Properties Examples

Size

Shape

Density

Color

Solubility

Texture

Melting Point

Chemical Properties

How matter reacts with other matter

Chemical Properties Example

Flammability

Corrosiveness

Acidity

Physical Change

Change in a physical property

Chemical composition unchanged

Easy to reverse

Chemical Changes

Change in chemical composition

Chemical Changes Example

Combustion

Photosynthesis

Physical Changes Example

Evaporation

Pure Substance

Constant characteristic physical and chemical properties

Elements

Pure Substances that cannot be broken down into smaller substances

Made up of one type of atom

Compounds

Substance made up of 2 or more different elements (atoms)

Different characteristics from the elements made up of

Diatomic elements (molecules)

2 or more atoms bonded together

H2

N2

O2

F2

Cl2

Br2

I2

Mixture

A combination of 2 or more substances

Can be present in varying amounts

Homogenous mixture

Uniform throughout

Looks like a single substance

Often be clear, light can go through unbothered

Heterogeneous Mixture

Not uniform throughout

Can see different substances

Measurements

a number and a unit plus a uncertainty

kilo (k)

1000

centi c

10^-2

mili (m)

10^-3

micro (μ)

10^-6

deci (d)

10^-1

Mega (M)

10^6

Giga (G)

10^9

Density

mass/volume or g/mL

Intensive property (independents of amount)

Extensive property

A property of matter that depends on the amount of substance present. Examples include mass and volume. It does not change with the size of the sample.

Intensive Property

Property that does not depend on the amount of substance present. It remains constant regardless of the quantity. Examples include temperature, density, and boiling point.

Volume

The amount of space occupied by an object or substance. It is measured in cubic units (e.g., cubic meters, cubic centimeters).

Volume: V = l × w × h (rectangular prism) V = πr²h (cylinder) V = 4/3πr³ (sphere) V = 1/3πr²h (cone) V = 1/2bh (triangle) V = πr² (circle)

Uncertainty in Measurement

There is always uncertainty, so often go one place past when measuring. This point is uncertain though, so you just add it. It refers to precision as well.

Significant Figures Rules:

• All non zeros are significant

• Zeros between non-zeros are significant

• Trailing zeros at end of number are only significant if a decimal point is present

• Leading zeros in front or before a number are never significant

  • Ex. 0.0024g has 2 sig figs

Adding and subtracting significant figures

The result should have the same number of decimal places as the term with the fewest decimal places

Multiplication and Division of Significant Figures

The result should have the same number of sig figs as the term with the fewest sig figs

Precision

• How close a series of measurements are to one another

  • Or how reproducible they are

Accuracy

How close a measured value is to the accurate (true) value

Exact numbers

• Infinite number of significant figures

  • Counting objects

  • Definition

    • 1,000 mL = 1 L

  • 2.54 cm=1in (exact)

Scientific Method

Observation, Hypothesis, and Experiments

Observation

Colleting data

Qualitative/quantitative

Hypothesis

• A tentative statement or interpretation of an observation

  • educated guess

• A good hypothesis is falsifiable

Experiments

Test hypothesis

Scientific Law

A brief statement on a series of observation which predict future ones

Law of Conservation of Mass

In a chemical reaction, matter is neither created nor destroyed

Scientific Theory

A model for the way nature is, and tries to explain what nature does, and why it does it

Predicts behavior fare beyond laws and observations

Law of Conservation of Mass

Mass of reactants equals mass of product

Law of Definite Proportions

All samples from a compound have the same proportion of constituent elements

Law of Multiple Proportions

When 2 elements form two different compounds, the masses of the elements that combine can be expressed as a ratio of small whole numbers

Dalton’s Atomic Theory

1. Each element is composed of tiny indestructible particles called atoms

2. All atoms of a given element have the same properties that distinguish them from atoms of other elements

3. Atoms combine in simple whole number ratios to form compounds

4. In chemical reactions, atoms are neither created nor destroyed. Nor can they be turned into other atoms

Problems with Dalton’s Atomic Theorey

• Sub-atomic particles smaller than atoms

• Nuclear reactions can turn one element into another element

• Isotopes- not all atoms are the same

Cathode Ray Tube Experiment

Experiment that demonstrated the existence of negatively charged particles (electrons) in atoms. Involved passing an electric current through a gas-filled tube with a cathode and anode. The cathode emitted a stream of electrons, which were attracted to the positively charged anode. Showed that the path of electrons could be manipulated by electric and magnetic fields.

Electrical Charge

• ± attract

• ++ repel

• - -repel

• Same repel, opposites attract

Millikan Oil Drop Experiment

What experiment measured the charge of an electron? It involved suspending oil droplets in an electric field and observing their motion.

The structure of an atom with the Plum-Pudding Model

Model proposed by J.J. Thomson in 1904. Atom consists of a positively charged "pudding" with negatively charged electrons embedded throughout.

Gold Foil Experiment

Experiment conducted by Ernest Rutherford in 1911 to probe the structure of atoms. Rutherford shot alpha particles at a thin gold foil. Most particles passed through, but some were deflected. This led to the discovery of the atomic nucleus and the idea that atoms are mostly empty space.

Proved the Plum-Pudding Model False

Nuclear model of the atom

The nuclear model of the atom states that atoms have a small, dense nucleus at the center, composed of protons and neutrons. Electrons orbit the nucleus in specific energy levels or shells.

Nucleus

• All positive

• Most mass of the atom

• Extremely dense

• Protons and neutrons

Electron Cloud

Mostly empty space

Atom is

electrically neutral

The number of electrons equals the number of protons

Proton

Mass of 1.0 amu

Charge +1

Neutron

Mass 1.0 amu

Charge 0

Electron

Mass is negligible

Charge -1

Elements

Defined by the number of protons

Atomic number

is the number of protons

identified by chemical symbol

Overall charge in elements =

protons - electroncs

Isotopes

Same number of protons

Different number of neutrons

Many different elements have different isotopes

Cations

Positive

Anion

Negative

Atoms can gain or lose electrons to

become cations or anions

Periodic Table of Elements (Left Side)

• Shiny

• Conduct heat and electricity

• Malleable

• Ductile

• Solids except Hg (mercury)

• Generally lose electrons in chemical reactions

Non-metals (Top right corner)

• Dull

• Don’t conduct

• Insulators

• Brittle

Metalloids (semi-metals)

Properties in between metals and nonmetals

Main Group of Periodic Table

1,2,13-18

Elements in the same vertical column

have the same properties

Alkali Metals (Column 1)

Violently Reactive

Alkaline earth metals (Column 2)

Less reactive than column 1

Noble Gasses (Colum 18)

Unreactive

Halogens (Row 17)

Reactive non-metals

All ions want to

be like the noble gasses, so gain and lose electrons

Atomic mass

Weighted average based upon natural abundance of each isotope

Atomic Mass=

Atomic Mass = (Fraction of isotope 1) + (mass isotope 1) + etc.

Natural abundance should all add to

100

Elements combine

to form a limitless number of compounds

Ionic Bonds (no lines)

• Electrons transferred between atoms (valence electrons)

• Bonds held together by electro-static forces

• Metals lose electrons to become cations

• Oppositely charge ions held together by ionic bonds, forming crystalline lattis

Covalent bonds

• Bonding atoms share electrons

• Electrons interact with booth nuclei

• Water is held together by covalent bonds

Molecular Formula

Actual number of atoms

Empirical Formula

Relative number

Simplified

Ex.: OH

Structural Formula

• Shows how atoms are connected

• Shows what bonds are present

Molecular Compound

2 or more covalent bonded nonmetals

Ionic Compound

Composed of cations and anions with a Formula Unit that is part of a lattice

Polyatomic ions

Ions composed of 2 or more atoms

Ionic Compounds: Formulas

• Ionic compounds contain positive and negative ions

• The overall charge must be neutral

• Formula reflects the smallest number ratio

How to formula Ionic Compounds

1. Find charge on metal cation, write it down

2. Find charge on nonmetal anion, write it down

3. Charge on cation becomes subscript on anion

   1. Charge on anion becomes subscript on cation

4. Lowest ratios? Neutral charge? Double check@

Name binary ionic compounds (which ones)

Metals only form one type of ion

Such as:

Li, Na, K, Rb, Cs

How to name binary compounds with metal (only form one):

Name of cation (metal) + base name of anion + -ide