The Structure of MatterPhysical Science

Chapter 6

Physical Science chapter 6 2

Review Compound: atoms of two or more elements

that are chemically combined Most of the matter around us is a compound or a

mixture of compounds Compounds have properties unlike those of their

elements During a chemical change, a new substance is

produced.

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Chemical bonds Forces that hold together the atoms in a

compound When atoms gain, lose, or share electrons

they are forming chemical bonds.

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Chemical Structure The way the atoms are bonded in a compound

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Chemical formulas Used as shorthand for writing compounds.

NaCl is sodium chloride Subscript – means written below

Tells us how many atoms of an element are in a compound

If there is no subscript, then there is one. Example: H2O has 2 atoms of hydrogen and one

atom of oxygen The ratio of hydrogen atoms to oxygen atoms is 2

to 1

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Chemical structure representations Chemical formula – show how many of each

type of atom there is Water: H2O Methane CH4

Structural formula – shows how atoms are arranged.

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Chemical structure representations Space filling model – shows relative volumes

of the electron clouds.

Ball-and-stick model – shows bond angles

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Effects of chemical bonds Compounds with strong chemical bonds

Are rigid and difficult to break Have high melting and boiling points

Compounds made of molecules Have strong bonds within each molecule Have weak attractions between molecules

Molecules are easy to separate Lower melting and boiling points

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Attractions between molecules Some molecules have stronger attractions

between them Example: water

Has hydrogen bonding between molecules Why water has a relatively high boiling point for

a molecular compound

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Discuss1. Classify the following

as mixtures or compoundsa. Air

b. CO

c. SnF2

d. Pure water

2. Draw a ball-and stick model of a boron trifluoride, BF3, molecule. A boron atom is attached to

three fluorine atoms. Each bond angle is 120 degrees and each bond is the same length.

3. Predict which molecules have a greater attraction for each other: C3H8O molecules in liquid rubbing alcohol or CH4 molecules in methane gas.

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Ionic compounds Ionic compound – a compound made up of

two or more ions Form networks of ions, not individual units. Ionic bond – the force that holds the ions in

an ionic compound together. Ionic compounds have a net charge of zero,

so the compound is electrically neutral.

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Examples NaCl

MgF2

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Ionic compounds Smallest unit is a formula unit. Generally have high melting points and high

boiling points. Are usually crystalline solids at room

temperature.

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Covalent compounds Covalent compounds are composed of

molecules that are created when atoms share electrons

Covalent bonds – the bonds between atoms in a molecule.

Molecules are also neutral.

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Examples HCl

Cl2

N2

O2

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Covalent compounds Smallest unit is a molecule. Generally have low melting points and

boiling points. Are usually liquid or gaseous at room

temperature, but not always.

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Discuss1. Determine if the

following compounds are likely to have ionic or covalent bonds.a. Magnesium oxide, MgO

b. Strontium Chloride, SrCl2

c. Ozone, O3

d. Methanol, CH3OH

2. Identify which two of the following substances will conduct electric current, and explain why.a. Aluminum foil

b. Sugar, C12H22O11 dissolved in water

c. Potassium hydroxide, KOH, dissolved in water

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Polar molecules Atoms in molecules don’t always share their

electrons equally.

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Examples HCl

H2O

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Polar molecule Has a positive end and a negative end. Example: stream of water

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Nonpolar molecules Do not have negative and positive ends. Example: CO2

Nonpolar vs. polar

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Metallic Bonds Occur between metal atoms Atoms are closely packed together Electron clouds overlap Electrons move freely between atoms

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Polyatomic ions in compounds Poly means many Polyatomic ions have more than one atom in

them. See Figure 10 on page 190

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Compounds with polyatomic ions They form compounds just like monatomic

(one atom) ions do. Examples

LiOH, lithium hyrdoxide

Mg(NO3)2, magnesium nitrate

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Discuss1. Compare bonds

2. Compare bonds

3. What is the difference between polar molecules and nonpolar molecules?

4. What are polyatomic ions?

5. Identify which of the bonds in calcium hydroxide, Ca(OH)2 are ionic and which are covalent.

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Naming Ionic Compounds List the (positive) cation first

Name is usually the same as the element List the (negative) anion second

Change ending to –ide See figure 2 on page 192

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Examples CaF2

Calcium fluoride

Li2O Lithium oxide

K2S Potassium sulfide

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Writing formulas for ionic compounds The charge on the compound must add up to

zero. Add subscripts as needed

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Examples Cesium Oxide

Cs2O

Beryllium chloride BeCl2

Calcium Phosphide Ca3P2

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Transition Metals Groups 3 – 12 Can have more than one charge when forming

compounds Copper and oxygen can make CuO or Cu2O3

To name them, we need to specify the charge of the cation using a roman numeral CuO is copper (II) oxide Cu2O3 is copper (III) oxide

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Examples Titanium (III) nitride

TiN

Fe2O3

Iron (III) oxide Iron (II) oxide

FeO

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Naming Covalent Compounds Different rules Use numerical prefixes (figure 5 on page 194) If there is only one atom of the first element, the

prefix mono- is omitted. Change the ending of the second element to -ide If the element starts with a vowel, drop the a or

o at the end of the prefix Example: tetroxide, not tetraoxide

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Examples PF5

Phosphorus pentafluoride

N2O5

Dinitrogen pentoxide

OF2

Oxygen difluoride Phosphorus trichloride

PCl3

Dinitrogen pentoxide N2O5

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Empirical Formulas Shows the smallest whole-number ratio of

atoms that are in a compound Ionic compounds

Almost always the same as the chemical formula Covalent compounds

Not always the same Example: glucose chemical formula is C6H12O6,

empirical formula is CH2O

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Molecular formula Shows the actual numbers of atoms of each

type in one molecule The same as the chemical formula

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Determining empirical formula Convert the mass of each element to moles. Find the molar ratio, which gives the

empirical formula The ratio must be whole numbers, because the

subscripts in the formula must be whole numbers. If the ratio isn’t whole numbers, multiply it by a

whole number to get rid of the fractions. Example: if the ratio is 1.5:1, multiply by 2 to get 3:2

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Examples A sample of an unknown compound has 36.04

g of carbon and 6.04 g of hydrogen. What is the compound’s empirical formula?

A sample of a compound contains 3.6 g of boron and 1.0 g of hydrogen. What is the compound’s empirical formula?

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You try A certain compound sample contains 207.2 g of

lead and 32.00 g of oxygen. What is its empirical formula?

A compound is analyzed and found to contain 36.70g potassium, 33.27g chlorine, and 30.03g oxygen. What is the empirical formula of the compound?

Find the empirical formula of a compound that contains 53.70g iron and 46.30g sulfur.