Chemistry 1A Final Exam Study Guide Name Hour

Chemistry 1A – Final Exam Study Guide Name Hour

Chapter 1 – Fundamentals of measurement and calculation

List the main metric prefixes and the meaning of each

For each number value given, underline the significant digits and list the correct number of significant digits  .00342

 140000

 12.010

 28.0 Convert each number given to proper scientific notation form  .0000048

 38700000 Convert each number given into proper decimal form  6.84x10-4

 5.2x106 Use dimensional analysis to convert between the metric units given:  Convert 12.9m to cm

 Convert .025 kg to g

 Convert .00kg to mg

State Standard C.1.7 o Define density and distinguish among materials based on densities. Perform calculations involving density.

A liquid has a mass of 32.1g and a volume of 43.8mL. Calculate the density.

Determine the mass (g) of a piece of wood with a density of .75 g/cm3 and a volume of 30.0 cm3

Determine the volume (mL) of a glass of water (density = 1.0 g/mL) that has a mass of 125g

Chapter 3 – States of Matter

State Standards o C.1.1 Based on physical properties, differentiate between pure substances and mixtures o C.1.2 Observe and describe chemical and physical properties of different types of matter o C.1.3 Recognize observable indicators of chemical changes o C.1.4 Describe physical and chemical changes at the molecular level o C.1.5 Describe the characteristics of solids, liquids and gases and changes in state at the molecular level Fill in the chart below to distinguish between solids, liquids and gases molecular definite shape? definite volume? fluid? spacing? solid liquid gas

Describe what is happening (at the molecular and macroscopic level) in each of the following physical changes  Evaporation

 Condensation

 Sublimation

 Freezing

 Melting

Explain the difference between a physical change and a chemical change (at the molecular level)

Compare and contrast elements, compounds, and mixtures at the molecular level and the macroscopic level

Explain the difference between a homogeneous mixture and a heterogeneous mixture. Identify which one is a “solution”

Explain how each of the following methods is used to separate a mixture  Settling

 Filtration

 Distillation

Chapter 4 – Atomic Structure and the Periodic Table

State Standards: o C.2.2 Describe how the subatomic particles (i.e., protons, neutrons and electrons) contribute to the structure of an atom and recognize that the particles within the nucleus are held together against the electrical repulsion of the protons. o C.2.3 Determine the number of protons, neutrons, and electrons in isotopes and in those isotopes that comprise a specific element. Relate these numbers to atomic number and mass number.

Fill in the following chart charge mass location

proton neutro n electro n

What is the significance of neutrons? Why are they necessary in an atom?

Use a periodic table to determine the following:  The number of protons in a silicon atom

 The number of neutrons in a bromine atom

 The most stable isotope of carbon

 The number of neutrons in a carbon 14 isotope

 The number of electrons in an oxygen atom

 The number of electrons in a chloride ion (Cl-1)

What is the difference between an atom and an ion?

How do ions form?

What type of elements usually form positive ions? Negative ions?

Write formulas for  Potassium nitride

 Magnesium bromide

 Calcium phosphide  Sodium iodide

In a formula for an ionic compound, which ion goes first? What does the suffix “ide” tell you? Chapter 5 – Formulas and Naming

State Standards C.3.4 Write chemical formulas for ionic compounds given their names and vice versa.

What does a suffix “ate” tell you? A compound between a metal and a nonmetal is usually A compound between 2 nonmetals is usually What category of elements often has more than one possible charge? How do you show the charge for those elements in the name of a compound? Write formulas for  Iron III chloride

 Iron III oxide

 Cobalt II sulfide

 Manganese IV oxide

Write formulas for  Silicon tetrachloride

 Dinitrogen pentoxide

 Boron triiodide

Write formulas for

 Lithium carbonate

 Potassium phosphate

 Calcium nitrate

 Ammonium sulfate

Write names for  CaBr2  CaCO3

 Li3P  Ba(NO3)2

 BF3  NiS

 SBR6  Cu2SO4

Chapter 11 – Electron Configurations

Indiana State Standards C.2.5 Write the electron configuration of an element and relate this to its position on the periodic table. C.2.6 Use the periodic table and electron configuration to determine an element's number of valence electrons and its chemical and physical properties.

Draw an outline of the periodic table and label the s, p, d and f blocks

Fill in the following tables

number of type of number of energy level orbitals type of orbitals orbital orbitals

Write the full electron configuration for  Aluminum

 Oxygen

Write the valence electron configuration for  Sulfur  Chromium

 Neon

Chapter 12 – Chemical bonding State Standards:

o C.3.1 Describe, compare and contrast the characteristics of the interactions between atoms in ionic and covalent compounds. o C.3.2 Compare and contrast how ionic and covalent compounds form. o C.3.3 Draw structural formulas for and name simple molecules. o C.3.5 Compare and contrast ionic, covalent, polar and non-polar molecular bonding

Explain the concept of electronegativity

What happens to the electronegativity value for elements as you  Go down a group (top to bottom)

 Go across a period (left to right)

Use the chart of electronegativity values on page 362. Determine the electronegativity difference to answer the questions  What kind of bond will form between C and Cl?

 What kind of bond will form between Mg and Cl?

 What kind of bond will form between H and O?

 Which is more polar, N to O or N to S?

 Which is more polar H to O or C to O?

Draw Lewis structures for  Phosphorus  HBr

 Aluminum  H2S

 Chlorine  PH3

 Argon  SiBr4

 Chapter 6

 Standard 4: Reactions and Stoichiometry  ------ Core Standard  Use balanced chemical equations and the mole concept to determine the quantities of reactants and products.  ------ C.4.1 Predict products of simple reactions such as synthesis, decomposition, single replacement and double replacement.  C.4.2 Balance chemical equations using the law of conservation of mass and use them to describe chemical reactions.  C.4.3 Given mass of the sample, use the mole concept to determine the number of moles and number of atoms or molecules in samples of elements and compounds.  C.4.4 Using a balanced chemical equation, calculate the quantities of reactants needed and products made in a chemical reaction that goes to completion.  C.4.5 Describe, classify and give examples of various kids of reactions-synthesis (i.e., combination), decomposition, single displacement, double displacement and combustion.  C.4.6 Determine oxidation states and identify the substances gaining and losing electrons in redox reactions.  C.4.7 Perform calculations to determine the composition of a compound or mixture when given the formula.