Chemistry

Answers to Explain (Analysis), Part 1:1. Examine the data collected for melting point. What conclusions can you draw about

the melting point of these chemicals?

- The chemicals that took longer to melt have a higher melting point than those that melted more quickly.

2. Which substances have higher melting points? Which have lower melting points? What does this indicate about the bonds in the substances?

  -Substances with ionic bonds had higher melting points and those with covalent bonds had lower melting points. A lower melting point indicates weaker bonds that will be more easily broken. A higher melting point indicates stronger bonds.

3. Summarize the solubility of the substances in the Explore Activity.

-All substances are soluble in water, except benzoic acid.

 

4. How is solubility associated with the type of bond present?

-All our ionic substances are soluble in water. However, not all ionic compounds are soluble in water. One of our covalent molecules are soluble (dextrose) and one is not (benzoic acid).

 5. What does the solubility of the different substances indicate about the type of bond present?

  -Substances with weaker intermolecular forces are more soluble in water. Those with stronger intermolecular forces are less soluble in water.

6. How is conductivity related to the type of bond present?

-Ionic substances, when dissolved in water, conduct electricity. Covalent substances do not conduct electricity when dissolved in water.

7. Why do substances with certain types of bonds conduct electricity well, while some substances are not good conductors?

-Ionic compounds conduct electricity well because they possess ions, which allows electrons to flow from atom to atom. This occurs only when they are melted or dissolved in water.

8. Is there a significant difference in appearance between the substances with covalent bonds and those with ionic bonds? What properties did you notice you could not see with the naked eye?

-Both ionic and covalent compounds appear to be white solids. Under the hand lens, however, the covalent substances are smaller in particle size than the particles in ionic compounds. Ionic particles also have a more geometric, crystalline shape while covalent particles vary in shape.

9. Imagine looking at the substances under a microscope. What do you think the substances might look like on a microscopic level?

-(Answers will vary.)

Lewis Structures are used to show bonding in molecules and ionic compounds.

-dots represent valence electrons -in ionic bonding, the charge of each ion must

be shown -in covalent bonding, bonded electrons is

shown by lines

Arrows represent transfer of electrons from the metal to the nonmetal.

-the charge of each atom must be shown

Example: CaS

When we show bonding, shared electron pairs can be shown by either a pair of dots or a single line.

-Lewis Structures are used to show how bonding electrons are arranged in molecules

-example: NH3

-sigma bond (): single covalent bond formed when

an electron pair is shared by the direct overlap of orbitals ♦can occur between s & s, s & p , or p & p orbitals

Explain(Analysis), Part 2:Draw Lewis structures for the following ionic

compounds. 1. NaCl 2. MgO 3. LiFDraw Lewis structures for the following

covalent compounds. 1. H2O

2. CO2

3. NH4

A multiple bond forms when two atoms share more than 2 electrons.

-double bond: 4 electrons shared ( 2 pairs) ♦ O2

-triple bond: 6 electrons shared (3 pairs) ♦ N2

Some molecules have both single and multiple bonds.

♦HCNpi bond (): forms when parallel orbitals overlap to

share electrons -only occurs with multiple bonds because the first overlap is always a sigma bond

Show the formation of the ionic compound for the following pairs of elements

1. strontium and fluorine2. aluminum and oxygen3. cesium and phosphorus4. lead and chlorine5. potassium and iodine6. magnesium and chloride7. aluminum and bromide8. cesium and nitride9. barium and sulfide

1. PH3

2. H2S

3. HCl4. SCl2

5. SiH4

6. CO2

7. CH2O

8. C2H2

chemical bond: force that holds two atoms together-creates stability in the atom

Two types of bonds:1. Attraction between a positive nucleus and

negative electrons (covalent bonding)2. Attraction between a positive ion and a negative

ion (ionic bonding)

Remember: It is the valence electrons that are involved in this bonding.

ionic bond: electrostatic force that holds oppositely charged particles together

-called ionic compounds -forms between metals and nonmetals ◊metals lose electrons, forms a cation

~cation: positive ion from loss of electrons ◊nonmetals gain electrons, forms an anion

~anion: negative ion formed from gain of electrons

-most are binary, which means they contain 2 different

elements, such as MgO, Al2O3

It is the chemical bonds between atoms that determines many of the physical properties of the compound.

-alternating positive and negative ions form an ionic crystal

-the ratio of positive to negative ions is determined

by the number of electrons transferred -strong attraction results in a crystal lattice, a 3-D arrangement of atoms.

Other characteristics include:-high melting and boiling points-very hard and rigid-brittle-electrolyte when dissolved in water (aqueous

solution)

During chemical reactions, energy is either absorbed (endothermic) or released (exothermic)

-the formation of ionic bonds is always exothermic

lattice energy: energy required to separate one mole of ions of an ionic compound

-the more negative the lattice energy, the stronger the bond

Lattice Energies of Some Ionic Compounds

Compound Lattice Energy Compound Lattice EnergyName (kJ/mol) Name (kJ/mol)

KI -632 KF -808

KBr -671 AgCl -910

RbF -774 NaF -910

NaI -682 LiF -1030

NaBr -732 SrCl2 -2142

NaCl -769 MgO -3795

Depends on:1. smaller ions -more negative value because the

attraction is stronger between the nucleus and valence electrons

2. larger the positive/negative charge, the more negative the lattice energy because the attraction is stronger when more electrons are lost/gained

Lattice Energyies of Some Ionic CompoundsCompound Lattice Energy Compound Lattice Energy

Name (kJ/mol) Name (kJ/mol)KI -632 KF -808

KBr -671 AgCl -910RbF -774 NaF -910NaI -682 LiF -1030

NaBr -732 SrCl2 -2142NaCl -769 MgO -3795

1. Draw the Lewis dot notation showing the bonding between beryllium and chlorine.

2. What determines the properties of an element?

3. What is a crystal lattice?4. List 5 characteristics of ionic compounds.5. What is the difference between endothermic

and exothermic? Which occurs in ionic reactions?

6. What is lattice energy?7. What does lattice energy depend on?8. Which substance has a stronger bond: NaCl

or NaBr? Why?

Remember that atoms bond to increase stability, which occurs when an atom gets a full outer shell of electrons.

-in ionic bonding, one atom loses electrons (metal)

and another gains electrons (nonmetal) to form oppositely charged ions with a full outer shell

However, sometimes there is not a transfer of electrons, but a sharing of electrons.

-covalent bond: attractive force between atoms due

to the sharing of valence electrons

Covalent bonds can form between: -2 or more nonmetal atoms -metalloids (especially the ones to the right of the metalloid line) and nonmetals

molecule: when two or more atoms bond covalently

Covalent bonds can have either single bonds or multiple bonds.

-single bonds: 2 shared electrons (1 pair) -multiple bonds: 4 or 6 electrons shared (2 pair= double or 3 pair = triple)

1. low melting and boiling points.2. many vaporize readily at room temperature3. relatively soft solids (but not all, some are

gases/liq.)4. can form weak crystal lattices5. do not conduct electricity when dissolved in

water

These properties are due as a result of differences in attractive forces

-attraction between atoms within a molecules is strong

-attraction between different molecules is weak ~called intermolecular forces or van der Walls

forces

Types of Intermolecular Forces (van der Walls forces)1. dispersion force (induced dipole)2. dipole-dipole force3. hydrogen bonding

dispersion force (induced dipole) -occurs between nonpolar molecules -very weakdipole-dipole force

-occurs between polar molecules -the more polar the molecule, the stronger the

forcehydrogen bonding

-strong intermolecular force between the hydrogen end of one dipole and a fluorine, oxygen or nitrogen atom on another molecule’s dipole

All bonds can be broken, though some more easily than others.

-due to the strength of the bond

What affects bond strength?bond length: distance that separates the bonded

nuclei -determined by the size of the atoms and how many electron pairs are shared ♦larger the atom, the longer the bond length, the weaker the bond ♦more shared electrons gives a shorter, stronger

bond

When a bond forms or breaks, an energy change occurs.

-bond formation: energy released (exergonic) -bond breaking: energy absorbed (endergonic)

bond dissociation energy: amount of energy required to break a specific covalent bond

-always a positive number -indicates the strength of a covalent bond

larger the bond dissociation energy, stronger the bond

(see p 246 for examples)

1. Describe a covalent bond.2. What types of atoms do covalent bonds form

between?3. Describe single, double and triple bonds.4. What do we mean by sigma and pi bonds?5. What do we call covalent compounds?6. What affects bond strength?7. Describe the two things that determine bond

length.8. What does bond dissociation energy indicate?9. What occurs when a bond forms or breaks?

Remember that atoms have different attractions for electrons (electronegativity).

-electronegativity increases left to right and decreases

down a period

The character and type of bond can be predicted using the difference in electronegativities between bonded atoms.

-pure covalent bond: electronegativity difference = 0

(usually occurs between identical atoms, H2)

Most atoms do not have equal sharing of electrons, producing a purely covalent bond.

-polar covalent bond: unequal sharing of electrons ♦the larger the electronegativity difference, the

more ionic the bond character -ionic bonds form when the electronegativity difference is > 1.7 and nonpolar covalent bonds

form when the difference is < 0.5 -the cutoff between polar covalent, nonpolar, and ionic is sometimes inconsistent with experimental data

Remember: bonding is not clearly ionic or covalent, with ionic character increasing as the difference in electronegativity increases.

Decide if the following pairs of atoms are polar covalent, nonpolar covalent or ionic.

1. N-H3.04-2.20 = 0.84

polar covalent2. C-Cl 2.55-3.16 = 0.61 polar covalent3. S-Se

2.58-2.55 = 0.03 nonpolar covalent

When a polar bond forms the shared electrons are pulled more strongly toward one atom.

-this creates partial charges at opposite ends of the molecule, which is called a dipole

♦ - indicates a partial negative + indicates a partial positive

Polar molecule or not?A molecule can have individual polar bonds, but

make a nonpolar molecule. How?We look at the shape of the molecule.

Let’s look at H2O and CCl4.

O—H C—Cl - + + -

1.24 0.61both O-H and C-Cl have polar covalent

bonds

One molecule is polar and the other is nonpolar? How do we know?

We look at the shape of the molecule and the terminal atoms.

-symmetric molecules like CCl4 are nonpolar because the polar bonds cancel each other out.

CCl4

-asymmetric molecules like H2O are polar because the polar bonds do not cancel each other out.

H2O

If water is polar, why will oil not dissolve in it?Oil must be nonpolar because

A substance is only soluble (dissolvable) when combined with a like molecule.

“Like Dissolves Like”

hydrophobic- “fear of water”hydrophilic- “likes water”

1. What is electronegativity and what does it predict?2. What is the difference between a nonpolar covalent

bond and a polar covalent bond?3. What is a dipole and what indicates them?4. Describe the electronegativity trend both across a period and down a group.5. Are the following bonds polar or nonpolar covalent? a. H-Br b. C-O c. S-C6. Describe the relationship between polarity and solubility.7. What do we mean by symmetric and asymmetric?

Final Bonding Questions:1.Draw a table comparing the properties of ionic and

covalent bonds. -leave room to add more properties (we will

discuss the table and add more to it)2.What is a general definition of a bond?3.What are the two types of bonds? Describe each.4.What is the octet rule?5.What do we mean by polar or nonpolar?6.What is electronegativity? How do we use this in

bonding?7.What are intermolecular forces?

1.

2. A bond is a force holding two atoms together to create stability in an atom

3. An ionic bond is an attraction of oppositely charged ions due to a transfer of electrons from a metal atom to a nonmetal atom. A covalent bond is the sharing of electrons between nonmetals or nonmetals and some metalloids.

Ionic Covalenthigh melting/boiling point low melting/boiling pointelectrolyte in water nonelectrolyte in watercrystal lattice structure some form weak crystal latticeshard, brittle solids gases, liquids, relatively soft solidsmost dissolve in water some dissolve in water many vaporize at room temp

4. The octet rule states that atoms are stable if they have a full valence shell of electrons. For most atoms, the number is 8, but the period 1 elements are stable with 2.

5. A covalent bond is polar if there is an unequal sharing of electrons due to the electronegativity difference between the atoms. It is nonpolar if there is an equal sharing of electrons.

6. Electronegativity is the attraction an atom has for electrons. The more electronegative the atom, the stronger the attraction. We use electronegativity to determine the polarity of molecules.

7. Intermolecular forces are the force that holds atoms together. They can be weak, allowing atoms to be pulled apart easily, or strong.

TEST #1

structural formula: uses letter symbols and bonds to show relative positions of atoms

-one of the most useful -can be predicted for many molecules by

drawing Lewis structures -H is always an end (terminal) atom, never a

central atom -less electronegative atom is the central atom

(nm or metalloid closest to the left of the PT-usually)

CH2O

1. Predict the location of the atomsC is least electronegative & farthest to left

on PT, therefore it is the central atom2. Find the total number of electrons available for bonding.

1 C-4, 2 H-2, 1 O-6 for a total of 12 valence e-

3. Determine the number of bonding pairs12 valence e- / 2 = 6 electron pairs

4. Place one bonding pair (single bond) between the

central atom and each terminal atom.

H C O H

5. Subtract the number of pairs you used in step 4 from

the number of bonding pairs determined in step 3.

6 – 3 used = 3 e- pairs left

5. Subtract the number of pairs you used in step 4 from

the number of bonding pairs determined in step 3.

-take the remaining electron pairs and place electron

pairs around the terminal atoms to satisfy the octet

ruleH C O H

6. If the central atom is not surrounded by 4 electron

pairs, it does not have an octet -convert one or two of the lone pairs on a

terminal atom to a double or triple bond between that

terminal atom and the central atom

H C O H

Practice:1. CH3Cl 2. NBr5

Writing structural formulas for polyatomic ions is the same with one exception:

-the total number of electrons may differ due to the

negative and positive charge. ♦negative charge, more electrons are present

SO4-2 add two electrons

♦positive charge, less electrons are presentNH4

+1 subtract one electron

Let’s look at CO3-2.

-when one or more valid Lewis structure can be

written for a molecule, resonance occurs -let’s look at another resonance molecule/ion:

NO3-1

-each molecule/ion that undergoes resonance behaves as if it only has one Lewis structure

Some molecules do not obey the octet rule.

Three reasons exist:1. when a small group of molecules have an odd

number of valence electrons: -NO2 for a total of 17 valance electrons-one

unpaired electron on N

2. Some form with fewer than eight, though this is relatively rare:

-B in BH3 is stable with six because it only has 3 valence electrons.

3. When the central atom has more than 8 electrons, which is referred to as an expanded octet.

-can occur in elements that are found in period three

or higher elements (because of the d orbitals). -P in PCl5

(1 s orbital, 3 p orbitals, and 1 d orbital)

1. SO3

2. N2O

3. SF6

4. ClF3

5. SiF4

6. PO4-3

7. BF3

8. SO3-2

1. What is a structural formula?2. Describe resonance.3. List three reasons for exceptions to the octet rule.4. Name the following: a. BH3 b. SO2 c. PO4

-3

5. Write formulas for the following: a. sulfur trioxide c. chlorous acid b. hydrosulfuric acid 6. Draw structural formulas a. SO2 b. H2O c. BrCl5

Many of the physical and chemical properties of molecules is determined by the shape of the molecule.

-the shape of molecules determines if two or more molecules can get close enough for a reaction to occur.

VSEPR (Valence Shell Electron Pair Repulsion) model: atoms in a molecule are arranged so that the pairs of electrons (bonded and lone) minimize repulsion.

The repulsion between electron pairs result in fixed angles between atoms

-bond angle: angle formed by any two terminal atoms and the central atom

♦lone pairs take up slightly more space than bonded

pairs ♦multiple bonds have no affect on the geometry because they exist in the same region as single bonds -example: H2O

See page 260 for the Molecular Geometries (Shapes)

1. What determines many of the physical and chemical properties of molecules?

2. Describe the VSEPR model.3. What does the repulsion between electron pairs result

in?4. Why do multiple bonds have no affect on geometry of a molecule?5. Why do molecules with lone pairs have shorter bond

angles?6. How do we know if a molecule is polar or nonpolar?

A universal set of rules must be used so chemists around the world can communicate.

formula unit: simplest ratio of ions represented in an ionic compound

-remember that ionic compounds form a crystal lattice, consisting of many cations and anions.

-the overall charge for the compound is 0

Most ionic compounds are binary, consisting of two monatomic ions.

-monatomic ion: one atom ion, either positively or negatively charged

Remember that we determine the charge of each ion by its oxidation number.

Formula Rules for Ionic Compounds1. write the cation first, followed by the anion 2. state the charges of both ions3. cross the number for the charge of one ion to

become the subscript for the other ion. -subscripts are used to state the number of

each atom in the compound

Example: Determine the formula for the ionic compound formed when potassium reacts with oxygen.

1. Cation = potassium = K Anion = oxygen = O2. K+1 O-2

3. K+1 O-2

K2O1

K2O

You try: Determine the formula for the ionic compound formed when aluminum reacts with chlorine.

We write formulas for ionic compounds containing polyatomic ions the same way as in binary compounds.

-the cation comes first, followed by the anion -state the charges -cross over the number for the chargesHowever: -if you have more than one polyatomic ion, place parenthesis around the polyatomic ion, with the subscript outside the parenthesis.

Example: Determine the formula for the ionic compound formed when beryllium reacts with cyanide.

1. Cation = beryllium = Be Anion = cyanide = CN-

2. Be+2 CN-1

3. Be+2 CN-1

Be1(CN)2

Be(CN)2

You try: Determine the formula for the ionic compound formed when ammonium reacts with iodine.

Write the correct formula for the following pairs of atoms:

1. ammonium and oxygen

2. lithium and nitrate

3. aluminum and hydroxide

4. ammonium and phosphate

5. strontium and acetate

1. Why do we need a universal set of rules for naming and writing formulas?

2. Define monatomic and binary.3. What is meant by a formula unit?4. What is the purpose of subscripts.5. Describe what a polyatomic ion is?6. When do we use parenthesis for writing ionic

compounds with polyatomic ions?7. Determine the formula for the ionic compound

formed when lead reacts with sulfur.8. Determine the formula for the ionic compound

formed when lithium reacts with nitrogen.

Write the correct formula for the following pairs of atoms:

1. aluminum and carbon2. ammonium and carbonate3. calcium and oxygen4. aluminum and chromate5. sodium and phosphate6. potassium and hydrogen sulfate7. magnesium and phosphorus

The names of ionic compounds include the ions of which they are composed.

1. The element whose symbol appears first in the formula also appears first in the name.

-this is always the positively charged ion, or metal

2. The name of the second ion follows, with its ending

changed to –ide for single atom ions.

Ex: What is the name of MgCl2?

magnesium chloride

Write the formula and the name.

1. Na2S

2. Ga2S3

3. CaSe

4. LiF

You follow the same rules when naming polyatomic ions as when you have binary ionic compounds, however:

-you do not change the ending of the polyatomic ions, even when they are the second atom.

Example: Al2(SO4)3

aluminum (III) sulfate

Rule: You must state the charge of all metals not included in groups 1 and 2 because many

have multiple charges.

Name the following compounds:1. NaC2H3O22. CaCO3

3. KOH4. Mg(NO3)2

5. Li2CrO4

6. Mg(NO2)2

7. AlPO4

*According to the previous rules, FeO and Fe2O3 would both be named iron oxide,even though they are not the same compound*

Since many transition metals can have more than one charge, the name must show this. This is done using roman numerals. -FeO is named iron (II) oxide because Fe has a +2

charge -Fe2O3 is named iron (III) oxide because Fe has

a +3 charge

*The roman numeral states the charge of the metal*

Q: How do I know the iron in FeO has a +2 charge? A: The oxide ion has a –2 charge, so the Fe must have a +2 charge so the compound is overall neutral.

Q: How do I know the iron in Fe2O3 has a +3 charge?

A: There are three oxide ions with a –2 charge:(3 ions)(-2 charge/ion) = a total of –6 charge

Since the overall charge must be neutral, the iron must have a total charge of +6. Therefore:

(2 ions)(x charge/ion) = +6 x = +3

Write the formula given & the name of each compound.

1. FeCl3

2. Zn3P2

3. CuS4. AuF5. CuC2H3O2

6. AgHCO3

7. ZnSO4

8. Pb(CO3)2

1. What is the ending of the second atom changed to when naming ionic compounds?

2. Write the name for (NH4)3P

3. Write the name for AlS.4. Determine the formula for the ionic compound

formed when magnesium reacts with phosphate.

5. Determine the formula for the ionic compound formed when magnesium reacts with phosphate.

Molecules are represented by both names and formulas.

Rules for Naming Binary Molecular Compounds1. The first element in the formula is named first, using the entire element name.2. The second element in the formula is named using the root of the element and adding the suffix –ide.3. Prefixes are used to indicate the number of atoms of each type that are present in the compound. -exception: 1st element never uses the prefix mono- -drop the final letter of the prefix if element name begins with a vowel.

Prefixes you need to know:# atoms prefix

1 mono- 2 di-

3 tri- 4 tetra-

5 penta- 6 hexa-

7 hepta- 8 octa- 9 nona- 10 deca-

Name the compound P2O5, which is used as a drying and dehydrating agent.

1st atom: P = phosphorus 2nd atom: O = oxygen = oxide

There are 2 phosphorus = diphosphorusThere are 5 oxygens = pentoxide (drop the –a of

penta-)

Put it together: diphosphorus pentoxide

Name the following molecules:1. CCl4

2. As2O3

3. CO

4. SO2

5. NF3

(We will talk more about acids in Ch 19)There are two types of acids:1. binary acid: contains hydrogen and one other element -when naming use the prefix hydro- plus the root

of the second element with the suffix –ic, followed

by the word acid. -ex: HCl

H = hydro- Cl = chloride = chloric hydrochloric acid

Some acids are not binary, but are named according to the binary acid rules when oxygen is not present, as in HCN.

H = hydro CN = cyanide = cyanic

hydrocyanic acid2. oxyacid: an acid that contains an oxyanion

(oxygen containing polyatomic ion) -the name depends on the oxyanion present

-the name consists of the root of the anion, a suffix,

and the word acid ♦if the anion suffix is –ate, it is replaced with -ic ♦if the anion suffix is –ite, it is replaced with -

ous

-examples: ~HNO3

NO3 = nitrate

= nitric nitric acid ~HNO2

NO2 = nitrite

= nitrous nitrous acid

Name the following acids:1. HBr

2. H3PO4

3. H2SO4

4. H2SO3

5. H2CO3

Use the prefixes in the molecule’s name to determine the subscript for each atom in the compound.

- phosphorus tribromide P Br

1 (no prefix) 3 (tri) PBr3

- the formula for an acid can be derived from the name as well

♦charge of the oxyanion or anion gives the number

of hydrogens hydrofluoric acid = HF

(1 H because fluorine has a -1 charge)

1. oxygen difluoride

2. dinitrogen tetrasulfide

3. phosphorus pentachloride

4. iodic acid

5. phosphoric acid

TEST # 2

Metallic bonds are similar to ionic bonds because they often form lattices in the solid state.

-eight to twelve metal atoms surround another, central metal atom

Instead of sharing electrons or losing electrons, the outer orbitals overlap.

-electron sea model: all metal atoms in a metallic

solid contribute their valence electrons to form a ‘sea’

of electrons around the metal atoms. -valence electrons are free to move from atom to atom (delocalized electrons), forming metallic cations

metallic bond: attraction of a metallic cation for the delocalized electrons that surround it

This bonding contributes to the unique properties of metals:

1. generally have high melting and boiling points, with

especially high boiling points -due to the amount of energy needed to separate the electrons from the group of cations 2. malleable (hammered into sheets) and 3. ductile (drawn into wire) -mobile electrons can easily be pulled and

pushed past each other

4. durable -though electrons move freely, they are strongly attracted to the metal cations and are not easily removed from the metal

5. good conductors -free movement of the delocalized electrons, allowing heat and electricity to move from

one place to another very quickly 6. luster -interaction between light and delocalized electrons

As the number of delocalized electrons increases, as in transition metals (d electrons), the hardness and strength also increases.

-alkali and alkaline earth metals are soft (s valence

electrons only)

It is easy to combine 2 or more different metals to make a metallic crystal

-alloy: mixture of elements with metallic properties

-the properties of alloys differ from those of the individual elements that make it up

1. What is a metallic bond?2. What is an alloy?3. Describe the electron sea model.4. What occurs with orbitals in metals?5. How is metallic bonding similar to ionic

bonding?6. What are delocalized electrons?7. What contributes to a metal’s high boiling

point, malleability, ductility and conductivity?8. List the other 2 properties of metals.9. What happens to strength and hardness as you

decrease the number of delocalized electrons?