Chapter 20: CHEMICAL BONDING

Chapter 20 Assignments• Tuesday 4/30 Ionic Bonding• Wednesday 5/1 Covalent Bonding• Thursday 5/2 Polar/Nonpolar Bonds• Friday 5/23 Oxidation Numbers• Monday 5/6 Writing Formula• Tuesday 5/7 Complex Ions• Wednesday 5/8 Writing Names• Thursday 5/9 Review- NTWS• Friday 5/10 Ch Review P626:1-18, 31-32• Monday 5/13 Catch up Review• Tuesday 5/14 Test

Classifications of Matter (p454)Matter

Yes Can it be separated by physical means?

Pure substance

Mixture

Can it be decomposed by ordinary chemical means?

Is the composition uniform?

ElementCompoundHeterogeneous

mixtureHomogeneous

mixture

No

Yes NoYes No

Yes No

Copper sulfate, CuSO4 is composed of

• 1 copper (Cu) atom• 1 sulfur (S) atom, and• 4 oxygen (O) atoms.

Chemical compounds (Page 602)• The atoms in compounds are chemically The atoms in compounds are chemically bondedbonded

and may only be broken down through chemical and may only be broken down through chemical reactions. reactions.

• A A chemical bond chemical bond is a mutual electrical attraction is a mutual electrical attraction between the nuclei and valence electrons of between the nuclei and valence electrons of different atoms. (p604)different atoms. (p604)

• Chemical bonds are electrical bondsChemical bonds are electrical bonds

Elements and Compounds

• Compounds formed when elements combine often have properties that are different from the properties of the elements from which the compound is formed.

• Copper Sulfate – greenish coating on the Statue of Liberty– Copper- shiny colored metal– Sulfur- yellow powder– Oxygen- clear, odorless gas

Sodium and Chlorine

• Na Sodium- a shiny soft, silvery metal that reacts violently with water.

• Cl Chlorine – a poisonous greenish-yellow gas

• NaCl Sodium Chloride- white crystals of ordinary table salt.

Chemical Formula (p603)

Familiar Name Chemical Name Chemical FormulaSand Silicon dioxide SiO2

Milk of Magnesia Magnesium hydroxide Mg(OH)2

Cane Sugar Sucrose C12H22O11

Lime Calcium Oxide CaOVinegar Acetic Acid CH3COOHLaughing Gas Dinitrogen oxide N2OGrain Alcohol Ethanol C2H5OHBattery Acid Sulfuric Acid H2SO4

Stomach Acid Hydrochloric Acid HCl

Why do elements combine to form compounds?

• Elements combine in ways that produce the most stable electron arrangement.

• The most stable electron arrangements have the outer energy shells filled.

• For elements with 1 or 2 electrons the filled outer shell has 2 electrons. Hydrogen and Helium (Helium is stable).

• For elements with > 2 electrons the filled outer shell has 8 electrons. Noble gases are all filled.

• This neutral lithium atom has three positively charged protons, three negatively charged electrons, and four neutral neutrons.

Element Structure

How do elements combine to form compounds?

• Atoms of elements gain, lose or share electrons in order to form a stable electron arrangement of 8 (or 2) electrons in the outer energy level.

• Whether an atom gains, loses or shares electrons depends on how strongly the atom attracts electrons.

• The closer a negatively charged electron is to the positively charged nucleus, the more strongly it is attracted to the nucleus.

• Moving down a group the outer level is farther from the nucleus

• The more protons in the nucleus of the atom, the stronger the attraction for the electron.

• Moving across a period the number of protons increases but distance stays the same.

• Close is more important than number of Protons.

How strongly does an atom attract electrons?

• Look at the horizontal rows, or periods, in the portion of the table shown.

Periodic Table and Energy Levels

Increasing attraction for electrons

Decreasing att

raction for electrons --

• You can determine the number of electrons in an atom by looking at the atomic number written above each element symbol. Group number allows us to determine number of outer electrons.

• Neon and the elements below it in Group 18 have eight electrons in their outer energy levels.

Noble Gases

• Their energy levels are stable, so they do not combine easily with other elements.

• The elements in Group 17 are called the halogens.

Halogens

• Fluorine is the most reactive of the halogens because its outer energy level is closest to the nucleus.

• Fluorine attracts electrons more than any other element.

Halogens• Note, that if Fluorine

gained an electron it would have the same electron arrangement as Neon.

Alkali Metals• Alkali metals each

have one outer energy level electron.

• It is this electron that is removed when alkali metals react.

• Unlike halogens, the reactivities of alkali metals increase down the group because the electrons are farther from the nucleus..

• The easier it is to remove an electron, the more reactive the atom is.

The Octet Rule• The octet rule states that chemical

compounds tend to form so that each atom, by gaining, losing, or sharing electrons has an octet, or 8 electrons, in its outer shell.

Bonding

Electrons are divided between inner shell and outershell electrons. Bonds form by the interaction of outer shell

or VALENCE ELECTRONS.

Valence ElectronsIA

IIA IIIA IVA VA VIA VII7A

VIIIA

For group 1 and 2 the number of valence electrons is equal to the group number. For group 13-18, the number of valence electrons is equal to the group number minus 10.

13 14 15 16 17

181

2

9 Ne

• Now that you know how to write electron dot diagrams for elements, you can use them to show how atoms bond with each other.

Using Dot Diagrams

• A chemical bond is the force that holds two atoms together.

• Atoms bond with other atoms in such a way that each atom becomes more stable. That is, their outer energy levels will resemble those of the noble gases. 8 electrons

Types of Chemical Bonds (P608)• There are two extreme forms of connecting or

bonding atoms:

• Ionic—complete transfer of electrons from one atom to another. One atom gains and the other loses.

• An atom that is no longer neutral because it has lost or gained an electron is called an ion (I ahn).

• Covalent—sharing of electrons between atoms

• Most bonds are somewhere in between.

Ionic Bonds—Loss and Gain (P609)(P609)• Sodium has only one electron in its outer level. • Removing this electron empties this level and

leaves the completed level below with 8 electrons which is a stable configuration.

• By removing one electron, sodium’s electron configuration becomes the same as that of the stable noble gas neon.

Na Na+

• Removing this electron also changes the sodium atom into a positively charged ion.

Ionic Bonds—Loss and Gain• Chlorine forms bonds in a way that is the

opposite of sodium—it gains an electron. • When chlorine accepts an electron, its electron

configuration becomes the same as that of the noble gas argon and the atom is changed into a negatively charged ion.

Cl Cl -

Bond Formation

• The positive sodium ion and the negative chloride ion are strongly attracted to each other.

• This attraction, which holds the ions close together, is a type of chemical bond called an ionic bond.

• A compound is a pure substance containing two or more elements that are chemically bonded.

Ionic BondsIonic BondsIonic BondsIonic BondsEssentially complete electron transfer

from an element of low attraction (metal) to an element of high electron attraction (nonmetal)

2 Na(s)+ Cl2(g) 2 Na+ + 2 Cl-

Therefore, ionic compounds exist primarily between metals at the left of the periodic table (Groups 1, 2, and transition metals) that form + ions and nonmetals at the right (Groups 16-17) that form - ions.

Na+Cl-

Lose 2 Gain 2• Some elements like

magnesium, Mg, in Group 2 has two electrons in its outer energy level.

• Magnesium can lose two electrons and achieve a completed energy level.

• Some atoms, such as oxygen (group 16), need to gain two electrons to achieve stability.

• The two electrons released by one magnesium atom could be gained by a single atom of oxygen. Magnesium Oxide is formed.

Lose 2 Gain 2 (group 2 and 16) Magnesium and Oxygen• When this happens, magnesium oxide

(MgO) is formed. • Mg Mg+2 +2e- and O+ 2e- O-2

• Mg+2 + O-2 ionic bond

• 2Mg +O2 2MgO

Magnesium and Chlorine

• Mg+2 + 2e- + 2 Cl Mg+2 +2Cl-1 MgCl2

Magnesium Chloride

Lose 2 Gain 1 (group 2 and 17)

Lose 1 Gain 2 (which groups)

Na+ O-2

Na2O sodium oxide

Metallic Bonding—Pooling• In a metal, the electrons in the outer energy levels of

the atoms are not held tightly to individual atoms. • Instead, they move

freely among all the ions in the metal, forming a shared pool of electrons.

• The outer electrons in metal atoms readily move from one atom to the next to transmit current.

Metallic Bonding—Pooling• Metallic bonds form when metal atoms

share their pooled electrons. This bonding affects the properties of metals.

Covalent Bonds—Sharing• Some atoms are unlikely to lose or gain

electrons because the number of electrons in their outer levels makes this difficult.

• The alternative is sharing electrons.• Many nonmetal elements form covalent bonds

with other non metals.

• These bonds follow the octet rule. • The chemical bond that forms between

nonmetal atoms when they share electrons is called a covalent (koh VAY luhnt) bond.

Olympic Day

• We need to pick a country-

• And we need to plan to win

• But-everybody participate!

The Covalent Bond• You can see how molecules form by

sharing electrons equally in this figure.

Electron Dot Diagrams of Electron Sharing

Cl Cl+ Cl Cl Cl2

Double and Triple Bonds• An atom can also form a covalent bond

by sharing more than two electrons. • When two pairs of electrons are shared by

two atoms, the bond is called a double bond.

Double and Triple Bonds• When three pairs of electrons are shared

by two atoms, the bond is called a triple bond. • Each nitrogen atom shares three electrons in forming a triple bond.

Forms of Chemical Bonds• There are two extreme forms of bonding atoms. Most

bonds are somewhere in between.• ionic • covalent

– Polar Covalent Bond — the electrons are not shared equally between atoms

– Read(P612-614)– Nonpolar Covalent Bond- – electrons are shared equally between

atoms

Diatomic GasesNon Polar Bonds (equal sharing)

• Br2- bromine

• I2 – iodine

• N2 – nitrogen

• Cl2 - chlorine

• H2 – hydrogen

• O2 – oxygen

• F2 – fluorine

• Mr. BrINClHOF

Polar Bonds and Polar Molecules• Some atoms of some elements have a

greater attraction for electrons than others do.

• A polar bond is a bond in which electrons are shared unequally.

Chapter 20 Section 3

Writing Formulas and Naming Compounds

Page 615

Chemical Formulas• A chemical formula is a combination of

chemical symbols and numbers (subscripts) that show which elements are present in a compound and how many atoms of each element are present.

• When no subscript is shown, the number of atoms is understood to be one.

• A water molecule contains one oxygen atom and two hydrogen atoms, so its formula is H2O.

• Binary Compounds- a compound composed of two elements.

Oxidation Number (P616)

• Oxidation number is the number of electrons lost, gained or shared by an element in a compound.

• The oxidation number can be determined from the periodic table.

• For Ionic compounds (ionic bonds) the oxidation number is the charge on the ion.

• Sometimes we are given the oxidation number (Table 2) for transition metals.

Oxidation Number (p616)Oxidation number is the number of electrons lost, gained or shared by an element in a compound. 1+ 2+ 3+ 4+ 3- 2- 1- 0

Special Ions (Table 2 P 616)Name Oxidation NumberCopper (I) 1+Copper (II) 2+Iron (II) 2+Iron (III) 3+Chromium (II) 2+Chromium (III) 3+Lead (II) 2+Lead (IV) 4+

Subscript/Superscript

• Subscript written below the symbol shows how many atoms or ions are present in the compound.

• Superscript written above the symbol shows the charge or oxidation number of the atom or ion and is used in knowing how to write the formula

• but is not actually written in the final formula.

1. Write the symbol of the element or polyatomic ion that has a positive oxidation number or charge. Hydrogen, the ammonium ion (NH4

1+), and all metals have positive oxidation numbers.

2. Write the symbol of the element or polyatomic ion with the negative oxidation number. Nonmetals other than hydrogen and polyatomic ions have negative oxidation numbers

3. The charge without the sign of one ion becomes the subscript of the other ion. Reduce the subscripts to the smallest common multiple.

• Mg2+ Cl1- MgCl2

Writing Chemical Formula (P617)

Step Magnesium Bromide Tin(IV)Sulfate

Write the symbols of the ions side by side with the positive ion first. Mg2+ Br1- Sn4+ SO4

2-

1. Cross over the charges by using the absolute value of each ion’s charge as the subscript for the other ion.

Mg2+1

Br1-2 Sn4+

2 (SO42-)4

1. Check the subscripts and divide them by their largest common factor to give the smallest possible whole-number ratio of ions.

Mg1 Br2 Sn2 (SO4)4

Write the formula Mg Br2 Sn(SO4)2

Example Writing Formula

• Al O Al O• Mg S Mg S• Li O Li O• K S K S• Ca N Ca N

Writing Names –Ionic Bonds(binary compounds) P618

1. Write the name of the positive ion.2. Some positive ions may have more than one

oxidation number. (see table 2) Use the Roman numeral to identify the ion charge. Negative ions have only one oxidation number.

3. Write the root name of the negative ion.Drop the last syllable

4. Add the ending –ide.

Determining the type of Bond

• Metallic Bonds- between transition metals and other transition metals

• Ionic Bonds- between elements on one side of the periodic table with an element on the other side. Non-metals bonding with metals.

• Covalent Bonds- between non-metals, elements close together on the right side of the PT.

Element Names in Binary Compounds

Element Name Name-ide oxygen oxidephosphorus phosphidenitrogen nitridesulfur sulfidefluorinebromine

Examples Writing Names

• BaF2

• NaCl• K2S

• MgCl2

Compounds with Complex Ions

• Binary Compounds – 2 elements• Complex Compounds

–Baking Soda NaHCO3

• Poly – many• Polyatomic- many atoms• Poly Atomic Ions- ions with more than 1

element.

Polyatomic Ions (p619)Charge Name Formula1+ Ammonium NH4

+

1- Acetate C2H3O2-

Chlorate ClO3-

Hydroxide OH-

Nitrate NO3-

2- Carbonate CO32-

Sulfate SO42-

Chromate CrO42-

3- phosphate PO43-

Naming Compounds with Polyatomic Ions

• First write the name of the positive ion– Either the metal or the ammonium ion

• Then write the name of the negative ion.• K2SO4

• ________ ______________• Sr(OH)2

• __________ _____________

Writing Formula From Name

• Same as for binary compounds but with the added caution that these ions must be in parentheses to show that the “subscript” goes with the entire ion.

Writing Formula with Polyatomic Ions1. Write the symbol of the element or polyatomic ion

that has a positive oxidation number or charge. Hydrogen, the ammonium ion (NH1+

4), and all metals have positive oxidation numbers.

2. Write the symbol of the element or polyatomic ion with the negative oxidation number. Nonmetals other than hydrogen and polyatomic ions have negative oxidation numbers

3. The charge without the sign of one ion becomes the subscript of the other ion. Reduce the subscripts to the smallest common multiple.

4. Put parentheses around the ions.

Practice Writing Formula

• Ammonium Phosphate• Write the positive ion and its charge• Write the negative ion and its charge• Balance the charges by X’ing - use ( ) .

Naming Covalent Binary Compounds• Naming Ionic Compounds - + element first, -

element second.• Covalent binary compounds are not formed

by ions bonding but with nonmetals sharing electrons.

• N2O, NO, NO2, N2O5 • We can’t rely on knowing charge to tell which

is which if we just say nitrogen oxide. • We use prefixes to tell how many of each

element.

Prefixes for Covalent Compounds

Number of Atoms Prefix1 Mono- Not often used2 Di-3 Tri-4 Tetra-5 Penta-6 Hexa-7 Hepta-8 Octa-

Examples

• N2O, • NO, • NO2, • N2O5 • CO2

• Exception CO

Hydrates• Ionic compound that has water chemically

attached to its ions and written into the chemical formula.

• cobalt chloride hexahydrate - pink crystals• CoCl2 * 6H2O

• Heat to drive out the water• CoCl2 anhydrous

• Plaster of Paris CaSO4

• add water forms calcium sulfate dihydrate• CaSO4 * 2H2O

Determining Oxidation Number

• Copper (I) or Copper (II)

• Copper(I) chloride or copper(II)chloride

• Cu1+ Cl1- Cu2+ Cl1-

• • Copper (I)Oxide or copper(II)oxide• Cu1+ O2- Cu2+ O2-

Put it all together• Given a formula

– Determine type of bond– Ionic bond: +name –name(ide)– Covalent bond: use prefixes to identify number of

atoms of each element

• Given a name– Determine type of bond– Ionic bond: determine oxidation numbers and

cross over to determine subscripts.– Covalent bond: use prefixes to determine

subscripts

Forms of Chemical Bonds• There are two extreme forms of connecting or

bonding atoms. Most bonds are somewhere in between.

• ionic – bonds between a metal and a non-metal• Electrons are traded

• Covalent- bonds between non-metals• Electrons are shared– Polar Covalent Bond — the electrons are not

shared equally between atoms– Nonpolar Covalent Bond— the electrons are

shared equally between atoms

Summary (p 602-606)

• When elements combine, the new compound has unique properties that are different from the properties of the original elements.

• Chemical symbols and numbers are shorthand for elements and their amounts in chemical formulas.

• The elements in group 18 (Noble Gases) have filled outer shells (either 2 or 8 electrons) and therefore rarely combine with other elements.

Summary (P608-610)• Electron dot diagrams show only the electrons

in the outer energy level (valence electrons).• Most atoms need 8 electrons to complete their

outer energy level.• Atoms form compounds by gaining, losing or

sharing electrons to complete their outer energy levels (8 electrons- octet rule)

• A chemical bond is the force that holds atoms together in compounds.

• The force is electrical.

Summary (p 610)• An ion is a charged particle that has either fewer

or more electrons than protons, resulting in a negative or positive charge.

• ion = an atom that has gained or lost 1 or more electrons

• An ionic bond is the force or attraction between opposite charges of ions in an ionic compound.

• An ionic compound is neutral because the sum of the charges of the ions is zero.

• Ionic compounds form lattice patters of multiple bonded ions.

Summary (p616)

• A binary compound is composed of two elements.

• The oxidation number tells how many electrons an atom has gained, lost or shared to become stable.

• The net charge (also net oxidation number) of a compound is zero.

Summary 611-614)• Some atoms, like those in Group 4 share

electrons instead of losing electrons or gaining them.

• Covalent bonds can form single or double or triple bonds.

• In polar molecules the electrons are shared unequally in the bond. This results in slightly charged ends.

• Electrons are shared equally in a nonpolar (pure covalent) molecule.

Summary (619)

• A polyatomic ion is a positively or negatively charged, covalently bonded group of atoms.

• A hydrate is an ionic compound that has water chemically attached to its ions.

• Greek prefixes are used to indicate how many atoms of each element are in a binary covalent compound.

Chapter 20 Schedule

• Monday- Practice, Covalent Compounds• Tuesday- Covalent Practice, Hydrates• Wednesday – practice and NTWS• Thursday – Chapter Review P626:1-18, 31-32

– Do not come to class with this incomplete (20pts)

• Friday- Chapter Test

Chapter 20 Review• Problem 31: What is the oxidation number of

Fe in the compound: You need to know this to write the name.

• Fe2 S3 Iron can be either 2+ or 3+

• Sulfur can only be 2-• Fe2 ? S3 2- total = 0

• 2(?) + 3(-2) = 0• ? = • Name:

Chapter Review• #32 Write the chemical formulas for the

following compounds.• A. Potassium chloride

• B. Calcium carbonate • • Copper (II) sulfate

• Sodium oxide

Put it all together• Given a formula

– Determine type of bond– Ionic bond: +name –name(ide)– Covalent bond: use prefixes to identify number of

atoms of each element

• Given a name– Determine type of bond– Ionic bond: determine oxidation numbers and

cross over to determine subscripts.– Covalent bond: use prefixes to determine

subscripts

Practice

Element Electrons in outer shell Electron dot diagram

cesiumcalciumarseniciodinebariumseleniumrubidium

PracticeName Formula Bond Type

lithium nitride

CuO

AlCl3

magnesium fluoride

(NH4)2O

Ca(OH)2

lithium nitrate

ammonium phosphate

Cr2O3

dinitrogen pentoxide

Na2 CO3

lead (II) oxide

Practice Problems• Write Formulas for the following compounds. Determine the

type of bond. For ionic compounds, find the charge number from the periodic table. Write the charge for each ion as a superscript. Cross over the charge superscripts to become the subscript of the other ion. For covalent compounds, use the prefix to determine how many atoms of the elements are present.

potassium iodide magnesium hydroxide aluminum sulfate chlorine heptoxideammonium carbonate

Practice Problems (page2)• Write the names of these compounds. Determine the type of bond

first. For ionic compounds, write the name of the positive ion first, then write the name of the negative ion. If the negative ion is an element, change the last syllable to “ide”. For covalent compounds use prefixes to tell how many atoms of each element are present.

KCl(chromium can have more than one oxidation state) Cr2O3

Ba(ClO3)2

NH4Cl

PCl3