General Chemistry

1st Semester Review

Chapter 2

• Qualitative observation: describes matter without using numbers– Ex:

• Quantitative observation: describes matter using measurements– Ex:

Classifying Matter by Composition

• A substance is matter with the same fixed composition and properties– Elements & compounds are substances

• A mixture is a combination of two or more substances in which the basic identity of each substance is not changed.– Can be separated by physical processes

Types of Mixtures

• A heterogeneous mixture does not have a uniform composition and its individual substances remain distinct.– Ex:

• A homogeneous mixture, or solution, always has a uniform composition and is the same throughout. – Ex:

– An alloy is a solid solution that contains different metals and sometimes nonmetallic substances.

Substances: Pure Matter

• There are two types of pure substances—compounds and elements.

Elements: the building blocks

• An element is the simplest form of matter– There are 117 elements – only 90 occur naturally

on earth– The periodic table organizes elements and uses

chemical symbols that are universally understood.

Compounds

• A compound is a chemical combination of two or more different elements joined together in a fixed proportion – Compounds have chemical formulas

Physical Properties

• Physical properties are those that do not involve changes in composition.

• Physical properties can be either quantitative or qualitative

When salt is dropped into water, the particles in the salt crystal

separate and are surrounded by water.

• Physical properties are characteristics of a sample of matter that can be observed or measured without any change to its identity.– Ex:

• A physical change is a change in matter that does not involve a change in the identity of the substance– Phase changes are physical changes!– Ex:

• Most matter on Earth exists in one of three physical states: solid, liquid, or gas.

• Changes in state are examples of physical changes because there is no change in the identity of the substance.

• Some substances are volitile, or change to a gas easily at room temperature.

Chemical properties and changes

• A chemical property can be observed only when there is a change in the composition of the substance

• Chemical change, otherwise known as a, is the change of one or more substances into other substances.

• According to the law of conservation of mass, matter is neither created nor destroyed in a chemical change

Energy & Chemical Changes

• Energy, which is the capacity to do work, is either absorbed or released during a chemical change– Energy has many different forms

Exothermic vs. Endothermic

• Exothermic reactions are chemical reactions that give off energy– Combustion

• Endothermic reactions are chemical reactions that absorb energy– photosynthesis

• Density is the amount of matter (mass) contained in a unit of volume.

• Density is a physical property

• Stays the same for a given substance– Doesn’t matter how much of how little you have

• Many possible units– g/m3, g/L, g/cm3, g/mL

• Density formula:

Sample Problem 1

• What is the density of a substance with a mass of 24.3 g and a volume of 32.9 mL?– Answer must have correct unit

Sample Problem 2

• What is the volume of an object with a density of 125 g/mL and a mass of 281 g?

• What is the mass of an object with a density of 4.36 g/mL and a volume of 500 mL?

Chapter 2 Atomic Theory

Dalton

1. all elements are composed of tiny particles called atoms which CANNOT be divided into smaller parts

2. All atoms of the same element have identical properties. Atoms of different elements have different properties

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

4. Chemical reactions take place when atoms rearrange. Atoms of one element are NOT changed into atoms of a different element

Thomson

• Discovered 1st subatomic particle!• Experiment

• Discovery

• Model

Rutherford

• Experiment

• Discovery

• Model

Particle Charge Location Mass

Proton

Neutron

Electron

Atomic Number & Mass Number

• Atomic number = number of protons• Differences among elements result from

different numbers of protons in their atoms

• Mass # = Protons + Neutrons• Electrons are so small they have almost no mass

• Atoms are electrically neutral• # of protons = # of electrons

Ions

• Atoms with electrical charge– Cation = • Forms by:

– Anion = • Forms by:

Isotopes

• Atoms with the same number of protons but different number of neutrons– Oxygen – 16, Oxygen – 17, Oxygen – 18– One isotope is more common than the others

Mass Number

• Distinguishes isotopes

• Protons + Neutrons = Mass Number– (Electrons have almost no mass)

Atomic Shorthand

• Beryllium- - -atomic number 4 mass number 9 m# p = 4 # e = 4 # n = 5

• The atomic number is written as a subscript.• The mass number is written as a superscript.

Be

• P =

• N =

• Electron =

– + ion– - ion

Element Atomic Number

# of Protons

# of Electrons

# of Neutrons

Mass #

2713 Al

2713 Al3+

Oxygen – 16

148O

9 10

20884 Po

53 54 74

Atomic Mass

• Different from Mass Number• Average of the masses of all the isotopes of

the element• Measured in Atomic Mass Units (amu)– Proton = 1 amu– Neutron = 1 amu– Electron = 0 amu

• Nitrogen has 2 naturally occurring isotopes, 14

7N has an abundance of 99.63%. 157N has an

abundance of 0.37%. What is the atomic mass of Nitrogen?

Chapter 3

Energy Levels

• Areas of space where electrons can move

• Closer to nucleus = lower energy• Further from nucleus = high

energy• ELECTRONS CANNOT EXIST

BETWEEN ENERGY LEVELS!!!• Numbered: level closest to

nucleus = 1

• S sublevel = 1 orbital• p sublevel = 3 orbitals• d sublevel = 5 orbitals• f sublevel = 7 orbitals• Each orbital can be filled: (2 e-), half filled: (1

e-), or empty: (0 e-)

Put it all together

• Energy level 1 = closest to nucleus– One sublevel = s– One orbital = 1s

• Energy level 2 – 2 sublevels = s and p– 4 orbitals = 2s, 2px, 2py, 2pz

• Energy level 3– Three sublevels = s, p, and d– Nine orbitals = 3s, 3px, 3py, 3pz, five 3d orbitals

• Energy level 4– Four sublevels = s, p, d, and f– 16 orbitals

• Higher energy levels have all four sublevels

Sample Problems

• What are the electron configurations of the following elements?

• Li

• O

• Cl

Sample Problems

• Co

• Kr

• Ba

Valence Electrons

• Electrons in the s & p orbitals of highest energy level

• Boron: 1s22s22p1

– Highest energy level– Valence electrons

• Scandium: 1s22s22p63s23p64s23d1

– Highest energy level– Valence Electrons

Dot Diagrams

• Valence electrons are the only ones that are involved in chemical reactions

• Dot diagrams show the valence electrons

1.Figure out how many valence electrons element has

2.Write the element symbol3.Add dots to top, right, bottom, and left of

symbol one at a time until all valence electrons are used

4.Remember: each orbital can only hold 2 dots

Element Electron Configuration

# of Valence Electrons

Dot Diagram

Phosphorus

Bromine

Oxygen

Chapter 4

Henry Moseley

• Rearranged the elements according to increasing atomic number– Atomic # =

• Resulted in the structure of the modern periodic table

• Periodic Law – the physical and chemical properties of the elements repeat in a regular pattern when they are arranged in order of increasing atomic number.

Periods and Groups

• 7 horizontal rows = periods– Correspond to outermost energy level

• Vertical columns = groups/families– Correspond to the number of outermost electrons– Have similar properties– Some have special names

Group/Family Names

• Group 1 = Alkali metals• Group 2 = Alkaline earth metals• Group 3 – 12 = Transition metals• Inner Transition metals• Group 17 (VII A) = Halogens• Group 18 (VIII A) = Nobel gases

Physical States and Classes of Elements

• Most elements are solid at room temperature

• Br & Hg are liquid

• N, O, F, Cl, and Noble gases are gas

• Elements are classified as metals, metalloids, or nonmetals based on their properties

• Metals– Located to the left of the stairs– Have luster– Conduct heat and electricity– Usually bend without breaking– Solid at room temperature– Very high melting points

• Nonmetals – located to the right of the stairs– Brittle– Dull looking– Poor conductors of heat and electricity– Usually gases

• Metalloids – have properties of both metals and nonmetals– B, Si, Ge, As, Sb, Te, Po, At– Some are semiconductors – conducts electricity

better than a nonmetal but not as good as a metal

Patterns in Valence Electrons

• Valence electron =

• All elements of a family have the same number of valence electrons

• Increase across a period

Why do atoms form ions?

• Representative elements lose or gain electrons in order to obtain the same electron configuration as a noble gas

Before

• Na– 1s22s22p63s1

• B– 1s22s22p1

• P– 1s22s22p63s23p5

• F– 1s22s22p5

After

• Na1+

– 1s22s22p6 (Ne)

• B3+

– 1s2 (He)

• P3-

– 1s22s22p63s23p6 (Ar)

• F1-

– 1s22s22p6 (Ne)

Bond Types

• Differences in electronegativity determine bond type

• 0 – 0.4 = nonpolar covalent• 0.4 – 1.7 = polar covalent• 1.7 – 3.3 = ionic

Nonpolar Covalent 0 – 0.4

• Electrons equally shared– Directly in the middle of the two atoms

• H2

• O2

Polar Covalent 0.4 – 1.7

• Electrons are shared by not equally– More electronegative atom pulls electrons closer

• HCl

• CCl4

Ionic >1.7

• Electrons not shared at all– Atom with higher electronegativity takes electrons

from atom with lower electronegativity• NaCl

• K2O

Molecule Electronegativity of 1st atom

Electronegativity of 2nd atom

Electronegativity difference

Bond Type

O2

HCl

NaCl

NO3

NH4

Rules for writing dot diagrams of molecules

1. Atoms want 8 electrons in their outer energy levels (octet rule)

2. Number of dots in the molecule = sum of dots in individual atoms

Examples

• CH4

• NH3

• CO2

VSEPR Theory

• V alence• S hell• E lectron• P air• R epulsion

• The shape of a molecule is determined by minimizing repulsion between lone pairs

Linear

• Atoms are in a line• CO2

Trigonal Planar

• Atoms make a flat triangle• CO3

2-

Tetrahedral

• Pyramid shaped• CH4

Trigonal Pyramidal

• A triangle that is not flat• NH3

Bent

• H2O

Naming Flow Chart

• Name the following:– KOH

– Ca(NO3)2

– (NH4)3PO4

– CuCl2

– CoN

– CCl4

– Hg3P2

Formula Writing Flow Chart

• Aluminum Sulfide

• Sodium Phosphate

• Calcium Nitrite

• Barium Oxide

• Carbon Tetrachloride

• Sulfur Hexaflouride

• Tetraphosphorus Decoxide

• Dinitrogen Monoxide

• Iron (III) Sulfate

• Titanium (III) Oxide

• Chromium (III) Carbonate

• Copper (I) Chloride

1000 m = 1km 1000 L = 1 kL 1000 g = 1 kg1 m = 10 dm 1 L = 10 dL 1 g = 10 dg1 m = 100 cm 1 L = 100 cL 1 g = 100 cg1 m = 1000 mm

1 L = 1000 mL 1 g = 1000 mg

1 m = 1000000 µm

1 L = 1000000 µL

1 g = 1000000 µg

Conversions – 1 step

Ex 1: A roll of wire is 15m long, what is the length in cm?

Ex 2: convert 8.96L to milliliters

1 step cont.

• Convert 100 yards to feet

• Convert 5 kilometers to miles

Conversions – 2 step

• The front board is 500 mm long, how long is it in km?

• A football field is 120 yards long, how long is it in miles?

2 step cont.

• Convert 525 km to cm

• Convert 10000 in to miles