Chapter 6 Bonding

Objectives Write the electron configuration of ions of representative elements. Define ionic bonds in terms of the difference in electronegativity of the atoms. Write the Lewis electron dot structures of atoms, ions and ionic compounds. Write the name of an ionic compound (binary, polyatomic and transition metal) given its formula, or its formula given its name. Describe the properties of ionic, metallic and covalent molecules and identify the forces holding them together.

Valence Electrons Mendeleev ordered his periodic table according to chemical behavior. A pattern he saw was that elements in the same group behaved the same chemically. This is due to its valence electrons. Valence electrons are the outermost electrons and are involved in bonding.

Calculating valence electrons Groups on the periodic table can help! What’s left out on the periodic table?

Transition Metals Transition metals can have different amounts of valence electrons depending on how they bond with nonmetals

Decide how many valence electrons each one has? Li F Kr Mg Group 13 Carbon Phosphorus 1 valence electron 7 valence electron 8 valence electron 2 valence electron 3 valence electron 4 valence electron 5 valence electron

Gilbert Lewis Lewis developed the idea of the octet and coined the term Lewis dot structures He was nominated 35 times for the Nobel Prize in chemistry, but never won.

Electron Dot structures A diagram that shows only the valence electrons around the atom as dots Some rules: Each side can only have 2 dots, for a maximum amount of 8 dots An Exception to group 18 is Helium, it will only have 2

Electron Dot structures Draw the correct Lewis dot structure for each given: 1 V.E 2 V.E 3 V.E 4 V.E 5 V.E 6 V.E 7 V.E 8 V.E

Try these on your own:

Electron dot structures

Think back to Nobel Gases Nobel gases are unreactive… What do you notice about their electron dot structures??

Octet Rule Atoms want to achieve stability or noble gas configuration by attaining 8 valence electrons. Maximum number in the s and p orbitals. Atoms do not want to gain or lose more than 3 electrons when bonding!

Which is the only group that has full octets? NOBLE GASES!

Ions

Ions Ions are formed when an atom loses or gains electrons Ions can gain or lose just one or multiple electrons to bond and achieve stability Ions do not want to gain or lose more than 3 electrons though when bonding

Cations Cations are formed when an atom loses electrons and becomes positive Metals form cations Examples: magnesium, potassium, aluminum

Anions Anions are formed when an atom gains electrons and becomes negative Nonmetals form anions Examples: fluorine, oxygen, phosphorus

How do we know if atoms lose or gain electrons? Let’s look at their Valence electrons Every atom wants to be at 8 valence electrons Sodium has 1 VE, so its easier to lose 1 and have a +1 Chlorine has 7 VE, so its easier to gain 1 and would have a – 1

Charge or Oxidation state The charge corresponds with the amount of valence electrons:

Did it lose or gain? Oxide ion Cesium ion Aluminum ion Bromide ion Phosphide ion O-2: gained 2 electrons Cs+1: lost 1 Al+3 lost 3 electrons Br-1 gained 1 electrons P-3 gained 3 electrons

Valence electrons: Gain or lose, and how many? Strontium Hydrogen Sulfur Xenon Iodine Lose 2 Lose 1 Gain 2 Neither, neutral Gain 1

Polyatomic Ions

Polyatomic Ions

Ions Review

Predicting Ionic Charges Group 1: Lose 1 electron to form 1+ ions H+ Li+ Na+ K+

Predicting Ionic Charges Group 2: Loses 2 electrons to form 2+ ions Be2+ Mg2+ Ca2+ Sr2+ Ba2+

Predicting Ionic Charges Group 13: B3+ Al3+ Ga3+ Loses 3 electrons to form 3+ ions

Predicting Ionic Charges Nitride Group 15: P3- Phosphide As3- Arsenide Gains 3 electrons to form 3- ions

Predicting Ionic Charges Oxide Group 16: S2- Sulfide Se2- Selenide Gains 2 electrons to form 2- ions

Predicting Ionic Charges F1- Fluoride Br1- Bromide Group 17: Cl1- Chloride I1- Iodide Gains 1 electron to form 1- ions

Predicting Ionic Charges Group 18: Stable Noble gases do not form ions!

Predicting Ionic Charges Groups 3 - 12: Many transition elements have more than one possible oxidation state. Iron(II) = Fe2+ Iron(III) = Fe3+

Predicting Ionic Charges Groups 3 - 12: Some transition elements have only one possible oxidation state. Zinc = Zn2+ Silver = Ag+

Chapter 6 Types of Bonding

Atoms seldom exist in nature as independent particles. Nearly all substances are made up of a combination of atoms that are held together by chemical bonds. Chemical Bond – a mutual electrical attraction between the nuclei and valence electrons of different atoms that bind the atoms together.

When atoms bond, their valence electrons are redistributed in ways that make the atoms more stable. The ways in which the electrons are redistributed determines the type of bonding. Metals tend to lose electrons to form positive ions, or cations, and nonmetals tend to gain electrons to form negative ions, or anions.

Chemical bonding that results from the electrical attraction between cations and anions is called ionic bonding. Chemical bonding that results from the sharing of electrons between atoms is covalent bonding. In covalent bonds the shared electrons are owned equally by the two bonded atoms.

You can estimate the type of bonding (ionic or covalent) between elements by the difference in electronegativities (Chapter 5).

Electronegativity 0.0-0.3 Non polar covalent 0.3-1.7 polar covalent This chart will help you determine if it is polar or nonpolar 0.0-0.3 Non polar covalent 0.3-1.7 polar covalent >/= 1.8 Ionic

Polar Covalent vs. Non-polar Covalent Blue shading represents electron density

Sample Problem Page 163 Use electronegativity differences and Figure 6-2 to classify bonding between sulfur, S, and the following elements: hydrogen, H; cesium, Cs; and chlorine, Cl as either ionic, polar covalent or non-polar covalent.

electronegativity difference bond type S and H 2.5 - 2.1 = 0.4 polar S and Cs 2.5 - 0.7 = 1.8 ionic S and Cl 3.0 - 2.5 = 0.5 polar

Ionic Bonds

Opposites attract! An ionic compound is composed of positive and negative ions that are combined so that the charges are equal. Cations will combine with anions to form ionic compounds or salts. (metal and nonmetal) This electrostatic force that holds them together is called an ionic bond!

Na+ + Cl- NaCl Example: sodium chloride consists of a positive ion (Na) with a +1 charge and negative ion (Cl) with a -1 charge. Na+ + Cl- NaCl Draw electron configurations and Lewis Dot Structures.

Ionic Bonds Cation + Anion  Ionic bond (Ionic compound) Salt (NaCl) Ionic bond between a metal and non-metal Look at the electron dot structures. Note how the sodium donates its electron to chlorine, now chlorine has an octet and a negative charge

Structure of Sodium Chloride Crystal

Structure of Sodium Chloride Crystal

Ionic bonds Show the ionic bonds of the following using the electron dot structure method. Magnesium and oxygen (MgO) 2. Potassium and Sulfur (K2S) 3. Calcium and chloride (CaCl2)

Properties of ionic compounds

Properties of ionic compounds Electrically neutral Hard, but brittle Most form crystal lattices at room temperature Generally, have high melting points and Boiling Points Cinnabar (HgS) Fluorite (CaF2)

Properties of ionic compounds Conduct electricity when dissolved in water or as a liquid.  Solids do not conduct electricity.

Ionic Compounds in Water When ionic compounds are placed in water, they will dissociate. This means the anions and cations will split apart and form weak bonds with the water molecules.

Sodium chloride vs sugar C12H22O11 NaCl Sodium is a metal Chloride is a nonmetal Melting point ~800oC C, H, and O are nonmetals Melting point ~185oC Which one (or both) is/are an ionic compound(s)?

Crystal Lattice Ionic compounds form in repeating patterns of anions and cations. Their crystal structure will be the same for a particular compound.

Ionic crystals Ionic compounds that are crystals are made out of small pieces called unit cells A unit cell is the simplest repeating unit in a crystal NaCl’s unit cell:

Lesson check: True or False: Ionic compounds have low melting points Nitrogen and oxygen form an ionic compound Ionic Compounds are normally liquids Ionic Compounds conduct electricity when dissolved False True

Metallic Bonding

What are some properties of metals? Think back to Chapter 1

Properties of metals Good conductors of electricity Ability to be drawn into a wire (Ductile) Ability to be hammered, without breaking (Malleable) Not brittle

Metallic bonds Metals are composed of closely packed cations held together by their outer electrons. These outer electrons can be referred to as a sea of electrons. The electrons move freely between metal atoms. This “sea of electrons” is what allows metals to be molded, hammered, or stretched.

Metallic bonding Metallic Bonds are the forces or attraction between those free floating outside electrons and the positively charged metal ions

Metal alloys An alloy is a mixture of 2 or more elements (one must be a metal) These are uniform throughout, so a homogeneous mixture Examples: Brass (copper and zinc); Sterling silver (silver and copper); Bronze (copper and tin)

Why have alloys? Alloys are important because they are combining properties and are often superior compared to the pure elements Typically, more inexpensive than the pure element: Sterling silver vs pure silver $0.95 vs $1.68

Think about this: A bronze statue is beginning to turn green; bronze is an alloy made of copper and tin Which element is causing it to become green? Hint: Think statue of liberty

Covalent Molecules

Molecules and Molecular Compounds Ionic bonds are a + and -, but what about CO2? What is type of elements are C and O? When nonmetals bond together a covalent bond is created and we call them molecules or molecular compounds!

Molecules Molecules are neutral atoms that are joined together by covalent bonds Molecular Compound another way a saying molecule Molecular formula shows you how many atoms of each element is in a substance Example: CO2 , NH4

Octet Rule and Covalent Bonding An octet is 8 valence electrons and want to achieve noble gas configuration! Molecules want the same thing, but they share their valence electrons

Sharing electrons Recall that ionic bonds give and take electrons… Molecules share their electrons between the 2 atoms. When they share their valence electrons, a covalent bond is made

Single covalent bonds When atoms share one pair of electrons they form a single covalent bond Example: H2 Let’s draw it:

Show these diatomics: Cl2 Br2 I2 F2 What about H2O?

Structural Formula Electron dot structure represents bonds as 2 dots coming together: A structural formula represents covalent bonds as dashes

What did we call those 2 dots next to one another? Lone pairs or unshared pairs They do NOT participate in bonding, but you must show them!

Try these on your own: NH3 CH4 H2O2 PCl3

Double Bonds Atoms that share two pairs of electrons Example: CO2

Double Bonds Draw O2:

Triple Bonds Atoms that share three pairs of electrons: Example: N2

Properties of Covalent Molecules

Properties of covalent molecules Made out of nonmetals Can be a solid, liquid, or gas at room temperature Low melting point and boiling points Poor to non-conductors of heat and electricity

H2O vs NaCl Liquid water Solid water

Strengths of covalent bonds vs. ionic bonds

Bonding Theories

How do we decide where to find electrons? The modern atomic theory tells us that they are most probable at certain locations

Molecular Orbitals When two atoms combine, their atomic orbitals combine and overlap to produce molecular orbitals A molecular orbital belongs to the whole of the molecule Each orbital can only contain 2 electrons When an orbital overlaps and participates in a covalent bond, it can be classified as a bonding orbital

Sigma bonding (σ) The first bond between sharing atoms is classified as sigma bonds

Pi bonding (π) The second type of bonding can be a pi Remember the first is a sigma and the rest can be pi bonds

Molecular Geometry VSEPR Theory

VSEPR Theory Valence-Shell-Electron-Pair-Repulsion theory This theory helps us understand the 3D structure of molecules and their properties. Bonding and unshared pairs of valence electrons become very important to us within VSEPR theory! The shapes of molecules are determined because electron pairs want to be far apart from each other (repulsion).

AXE – Method to represent compounds A represents the central atom X represents the bonding atoms E represents the lone pairs the central atom has

Compounds with no lone pairs

Bonding pairs are far apart from each other Draw or build CO2 Meaning 1 central atom, 2 bonded atoms It has a linear shape No lone pairs A bond angle of 180o Bonding pairs are far apart from each other

This has a trigonal planar Draw or build BF3 This has a trigonal planar Meaning 1 central, 3 bonded Bond angle: 120o Bonds pointing to the corners of a triangle

Draw or build CH4 This has a tetrahedral AX4 Bond angle: 109.5o Meaning 1 central, 4 bonded Bond angle: 109.5o

Problem Use VSEPR theory to predict the shape of: aluminum trichloride, AlCl3 hydrogen iodide, HI carbon tetrabromide, CBr4 dichloromethane, CH2Cl2

Compounds with lone pairs Lone pairs occupy space, but only bonded atoms determine the name

Draw or build H2O This has a bent shape AX2E2 Meaning 1 central, 2 bonded, 2 lone pairs Bond angle: 105o Similar angles to the tetrahedral bond angles

This has a trigonal pyramidal AX3E Draw or build NH3 This has a trigonal pyramidal AX3E Meaning 1 central, 3 bonded, 1 lone pair Bond angle: 107o Similar angles to the tetrahedral bond angles

What shape is each one? Linear BeCl2 Bent OF2 Trigonal Planar AlCl3 Trigonal pyramidal Tetrahedral BeCl2 OF2 AlCl3 PCl3 CF4

Switch presentations – slide 80 Bond Polarity Switch presentations – slide 80

Bond polarity Since, atoms are sharing within a covalent bond… If they share equally they are a nonpolar covalent bond Examples: Diatomic atoms are nonpolar because they pull on each other evenly

Bond polarity Since, atoms are sharing within a covalent bond… If they share unequally they are a polar covalent bond Examples: HCl, H2O

Bond polarity Since, polar bonds are unequally sharing we will have dipoles. But how will we decide polarity?? Electronegativity!!

Dipoles The more electronegative will have the arrow point towards it and have a slightly negative charge The less electronegative will have a slightly positive charge

Electronegativity 0.0-0.4 Non polar covalent This chart will help you determine if it is polar or nonpolar 0.0-0.4 Non polar covalent 0.4-1.0 Slightly polar covalent 1.0-2.0 Very polar covalent >/= 2.0 Ionic

Decide the polarity of the following: N-H F-F Ca- Cl Al- Cl 0.9 slightly polar 0 Nonpolar 2.0 ionic 1.5 very polar