2. UNIT 2: Structure and Properties of Matter
Chapter 3: Atomic Models and Properties of Atoms
Chapter 4: Chemical Bonding and Properties of Matter

3. Chapter 4: Chemical Bonding and Properties of Matter
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Chapter 4: Chemical Bonding
and Properties of Matter
The chemical bonding in a substance influences the shape of its molecules, and molecular shape influences the properties of that substance. One of the properties of iron is its strength, which makes it ideal for use in support structures.
The strength of iron makes it useful in items such as horseshoes.
TO PREVIOUS SLIDE

4. 4.1 Models of Chemical Bonding
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
4.1 Models of Chemical Bonding
Three types of chemical bonding are ionic, covalent, and metallic.
TO PREVIOUS SLIDE

5. Electronegativity UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Electronegativity
Electronegativity is the relative ability of an atom to attract shared electrons in a chemical bond.
Answer to question prompt: electronegativity tends to increase from left to right and decrease from top to bottom
What general trends in electronegativity are shown in the periodic table?
TO PREVIOUS SLIDE

6. Electron Sharing and Electronegativity
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Electron Sharing and Electronegativity
Electronegativity difference, ΔEN, between two atoms bonded together can be low, intermediate, or high. The electron density diagrams below show the differences in the bonds.
when ΔEN is 0: electrons are equally shared
when ΔEN is 1: electrons are more closely associated with the more electronegative atom
when ΔEN is high, there is little sharing of electrons
Bonding is a continuum between equal sharing and minimal sharing of electrons.
TO PREVIOUS SLIDE

7. Electron Sharing and Electronegativity
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Electron Sharing and Electronegativity
Scientists have categorized types of bonds according to ΔEN.
ΔEN between 1.7 and 3.3: mostly ionic
ΔEN between 0.4 and 1.7: polar covalent
ΔEN between 0.0 and 0.4: mostly covalent (non-polar)
Three categories of bonds have been set based on ΔEN .
TO PREVIOUS SLIDE

8. UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Metallic Bonding
Chemists use the electron-sea model to describe metallic bonding. The model proposes that the valence electrons of metal atoms move freely among the ions, forming a “sea” of delocalized electrons that hold the metal ions rigidly in place.
Microscopic analysis shows that the structure of metals consists of aggregates of crystals.
TO PREVIOUS SLIDE

9. Properties of Metals UNIT 2 Melting and Boiling Points
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Properties of Metals
Melting and Boiling Points
the stronger the bonding forces, the higher the melting and boiling points of pure metals
In Group 1, melting points decrease as the atomic number gets larger because the atoms in each period have one more electron shell than atoms in the previous period. Thus free valence electrons are progressively farther from the nucleus, and the strength of the attractive forces decreases.
Across a period, the melting points increase as the atomic number becomes larger because the number of valence electrons increase across a period for the first several groups. Therefore the ions have a larger positive charge and a stronger attractive force.
Periodic table trends include:
For Group 1, melting points decrease as the atomic number increases.
For Groups 1 to 6, across a period, melting points increase as atomic number increases.
TO PREVIOUS SLIDE

10. Properties of Metals UNIT 2 Electrical and Thermal Conductivity
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Properties of Metals
Electrical and Thermal Conductivity
Metals are good conductors because their electrons are free to move from one atom to the next.
Malleability and Ductility
Based on the electron-sea model, metals can be shaped because, when struck, the metal ions can slide by one another while the electrons still surround them.
Hardness
The variation between metals is due to differences in crystal size (smaller ones make harder metals).
TO PREVIOUS SLIDE

11. Alloys UNIT 2 Alloys are solid mixtures of two or more metals.
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Alloys
Alloys are solid mixtures of two or more metals.
the addition of the second metal, even in a very small amount, can significantly affect the properties of a substance
in some cases, non-metal atoms, such as carbon, are added
If atoms of the second metal are similar in size to the first metal, they take the place of those atoms.
If atoms of the second metal are much smaller than atoms of the first metal, they will fit in spaces between the larger atoms.
TO PREVIOUS SLIDE

12. Ionic Bonding UNIT 2 occurs when ΔEN is between 1.7 and 3.3
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Ionic Bonding
occurs when ΔEN is between 1.7 and 3.3
essentially, involves one atom losing one or more electrons and another atom gaining those electron(s)
There are different ways to show the transfer of electrons in the formation of ionic compounds.
TO PREVIOUS SLIDE

13. UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Ionic Crystals
Ionic compounds exist as crystal lattice structures with particular patterns of alternating positive and negative ions. The unit cell is the smallest group of ions that is repeated.
NaCl forms a cubic crystal lattice structure.
Different types of crystal structures can form.
the relative sizes and charges of the ions affect the type of crystal structure that an ionic compound will form.
TO PREVIOUS SLIDE

14. Properties of Ionic Compounds
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Properties of Ionic Compounds
Melting and Boiling Points
high due to very strong attractions between ions
Solubility
ionic compounds are soluble in water when the attractive forces between the ions and water molecules are stronger than the attractive forces among the ions themselves
When sodium chloride (NaCl) dissolves in water, attractive forces between water molecules and NaCl ions act to break apart the ionic bonds.
TO PREVIOUS SLIDE

15. Properties of Ionic Compounds
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Properties of Ionic Compounds
Mechanical Properties
hard and brittle, so will break apart when struck
Ionic crystal will break on smooth planes, where like charges become aligned.
Conductivity
solids do not conduct because ions cannot move
compounds conduct when dissolved in water and ions can move
TO PREVIOUS SLIDE

16. Covalent Bonding UNIT 2 occurs when ΔEN is less than 1.7
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Covalent Bonding
occurs when ΔEN is less than 1.7
covalent bonds are classified into two types:
polar covalent: atoms do not share electrons equally
non-polar covalent: atoms share electrons almost equally
Forces in covalent bonds:
both attractive and repulsive forces play a role
The length of a covalent bond is determined by different electrostatic forces.
TO PREVIOUS SLIDE

17. Answer on the next slide
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Learning Check
Describe the chemical bonding and structure of NaCl. How do bonding and structure influence the general properties of the substance?
Answer on the next slide
TO PREVIOUS SLIDE

18. UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Learning Check
NaCl is composed of a metal atom bonded to a non-metal atom with ΔEN > 1.7. As such, the bond is classified as ionic. It exists as a cubic crystal lattice structure, with an alternating pattern of chloride ions and sodium ions.
Properties of NaCl include high melting and boiling points; solubility in water; hard and brittle; a poor conductor as a solid, but it does conduct electricity when dissolved in water.
TO PREVIOUS SLIDE

19. Quantum Mechanics and Bonding
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Quantum Mechanics and Bonding
Quantum mechanics is used to explain and describe chemical bonding. It is also used to account for shapes of molecules.
Valence Bond (VB) Theory explains bond formation and molecular shapes based on orbital overlap.
The region of overlap has a maximum capacity of two electrons, which have opposite spins.
There should be maximum overlap of orbitals, since the greater the overlap, the stronger and more stable the bond.
Atomic orbital hybridization is used to help explain the shapes of some molecules.
The first principle of VB theory is in accordance with the Pauli exclusion principle.
The third principle, atomic orbital hybridization, involves the process of combining or mixing of atomic orbitals to produce new atomic orbitals.
TO PREVIOUS SLIDE

20. Quantum Mechanics and Bonding
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Quantum Mechanics and Bonding
Molecular Orbital (MO) Theory explains bond formation and molecular shapes based on the formation of new molecular orbitals.
According to MO theory:
Covalent bond formation involves atomic orbital overlap that results in formation of new orbitals called molecular orbitals.
Molecular orbitals have shapes and energy levels that are different from those of atomic orbitals.
The electrons in molecular orbitals are delocalized throughout the orbital.
TO PREVIOUS SLIDE

21. Explaining Single Bonds
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Explaining Single Bonds
For molecules like hydrogen fluoride:
the 1s orbital of H overlaps with the half-filled 2p orbital of F
According to MO theory, the bond is a sigma (σ) bond, which is symmetrical and freely rotates.
TO PREVIOUS SLIDE

22. Explaining Single Bonds
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Explaining Single Bonds
For molecules like methane:
the VB theory of hybrid orbitals is used to explain molecular shape
carbon forms four hybrid orbitals (sp3) by combining three 2p orbitals and a 2s orbital so that four identical bonds can be created
The four sp3 orbitals of C overlap with the s orbitals of H to form methane.
TO PREVIOUS SLIDE

23. Explaining Double Bonds
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Explaining Double Bonds
Hybrid orbitals are used to explain the structure of ethene or molecules like ethene.
it is planar with ~120º bond angles
the structure is explained by formation of 3 sp2 hybrid orbitals for each carbon (a 2s orbital mixes with two 2p orbitals)
TO PREVIOUS SLIDE

24. Explaining Double Bonds
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Explaining Double Bonds
For bond formation in ethene:
one sp2 orbital of each carbon overlaps to form a σ bond between the carbons
two sp2 orbitals of each carbon overlap with the 1s orbitals of the hydrogens to form σ bonds
the lobes of the 2p orbitals of each carbon overlap above and below the plane to form a pi (π) bond
TO PREVIOUS SLIDE

25. Explaining Triple Bonds
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Explaining Triple Bonds
For molecules like ethyne:
the linear structure is explained by formation of 2 sp hybrid orbitals for each carbon (a 2s orbital + a 2p orbital)
sigma bonds form from overlap between sp of each carbon and between sp of carbons and 1s of hydrogens
two pi bonds form from overlap of the two 2p orbitals of each carbon
TO PREVIOUS SLIDE

26. Types of Hybridization
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Types of Hybridization
The names of hybrid orbitals (formed by the combination of two or more orbitals in the valance shell of an atom) indicate the number and types of atomic orbitals that were combined.
Atoms of Period 3 elements can have d orbital hybridization with s and p orbitals.
The number of hybrid orbitals that form is the same as the number of atomic orbitals that are combined.
Examples of hybrid orbitals and how they influence the shape of a molecule is discussed in more detail in section 4.2.
Each hybrid orbital has a certain overall shape.
TO PREVIOUS SLIDE

27. UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Allotropes
Allotropes are compounds that consist of the same element but have different physical properties.
An example is allotropes of carbon, which differ in the pattern of covalent bonds between carbon atoms.
Allotropes of carbon: A graphite, B diamond,
C buckyballs, D nanotubes
TO PREVIOUS SLIDE

28. Covalent Network Solids
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Covalent Network Solids
Network solids are substances that consist of atoms bonded covalently in a continuous two- or three-dimensional array. There is no natural beginning or end to the chains of atoms.
Silicon dioxide, SiO2, exists as a network solid that is represented as (SiO2)n.
TO PREVIOUS SLIDE

29. UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.1
Section 4.1 Review
TO PREVIOUS SLIDE

30. 4.2 Shapes, Intermolecular Forces, and Properties of Molecules
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
4.2 Shapes, Intermolecular Forces, and Properties of Molecules
Molecular compounds form a much greater variety of structures than ionic compounds form.
Understanding the properties of molecules requires an understanding of their three-dimensional shapes.
Different theories and models are used to predict molecular shapes.
The shape of a molecule is the result of the presence of atoms, bonding electrons, and non-bonding electrons, as well as forces of attraction and repulsion.
TO PREVIOUS SLIDE

31. UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
Depicting Two-Dimensional Structures of Molecules with Lewis Structures
Drawing the two-dimensional Lewis structure is the first step to predicting the three-dimensional structure of a molecule.
TO PREVIOUS SLIDE

32. Some Exceptions When Drawing Lewis Structures
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
Some Exceptions When Drawing Lewis Structures
Co-ordinate Covalent Bonds:
one atom contributes both electrons
bonds behave the same way as other covalent bonds and therefore are not indicated in Lewis structures
The ammonium ion has a co-ordinate covalent bond.
Expanded Octet (Expanded Valence):
central atom has more than an octet of electrons
a feature of some Period 3 and higher elements
Some exceptions to the octet rule and standard approach to drawing Lewis structures are required for certain molecules and ions.
For SF6(g), 12 electrons are around the central atom.
TO PREVIOUS SLIDE

33. Some Exceptions When Drawing Lewis Structures
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
Some Exceptions When Drawing Lewis Structures
An Incomplete Octet:
central atom has fewer than an octet of electrons
In BF3(g), boron has an incomplete octet.
Resonance Structures:
measured bond lengths may not support Lewis structures
one of two or more Lewis structures that show same relative position of atoms but different positions of electron pairs
Actual bond lengths in ozone are between those of single and double bonds.
TO PREVIOUS SLIDE

34. Predicting the Shapes of Molecules Using VSEPR Theory
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
Predicting the Shapes of Molecules Using VSEPR Theory
The valence-shell electron pair repulsion (VSEPR) theory
is a model used to predict molecular shape
is based on electron groups around a central atom being positioned as far apart as possible (repulsion)
predicts certain arrangements of electron groups
An electron group is a single bond, double bond, triple bond, or lone pair of electrons.
A bond angle is the angle between the nuclei of two atoms that surround the central atom of a molecule.
Note: It is important to distinguish between electron-group arrangement (how electron pairs are positioned around a certain atom) and the relative positions of atoms in the entire molecule.
For VSEPR, there are five electron-group arrangements. (Electron groups are represented by bars).
TO PREVIOUS SLIDE

35. Electron Groups and Molecular Shapes
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
Electron Groups and Molecular Shapes
When all the electron groups around a central atom are bonding electrons (i.e., there are no lone pairs of electrons), then the molecular shape has the same name as the electron-group arrangement.
TO PREVIOUS SLIDE

36. Electron Groups and Molecular Shapes
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
Electron Groups and Molecular Shapes
If one or more electron groups around a central atom is a lone pair, different strengths of repulsive forces will alter bond angles to differing degrees.
If one or more electron groups around central atom is a lone pair, the molecular shape name differs from the electron-group arrangement name.
TO PREVIOUS SLIDE

37. Summarizing Molecular Shapes
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
Summarizing Molecular Shapes
A summary of the VSEPR-based electron-group arrangements and associated molecular shapes is shown on this slide and the following slide.
TO PREVIOUS SLIDE

38. Summarizing Molecular Shapes
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
Summarizing Molecular Shapes
TO PREVIOUS SLIDE

39. Guidelines for Using VSEPR Theory to Predict Molecular Shape
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
Guidelines for Using VSEPR Theory to Predict Molecular Shape
TO PREVIOUS SLIDE

40. Answer on the next slide
UNIT 2
Chapter 3: Atomic Models and Properties of Atoms
Section 3.2
Learning Check
What is the electron-group arrangement and molecular shape of HCN?
Answer on the next slide
TO PREVIOUS SLIDE

41. Learning Check UNIT 2 HCN has two bonding groups and no lone pairs.
Chapter 3: Atomic Models and Properties of Atoms
Section 3.2
Learning Check
HCN has two bonding groups and no lone pairs.
The electron-group arrangement is linear, and the shape of the molecule is also linear.
TO PREVIOUS SLIDE

42. Determining the Hybridization of the Central Atom of a Molecule or Ion
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
Determining the Hybridization of
the Central Atom of a Molecule or Ion
TO PREVIOUS SLIDE

43. Answer on the next slide
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
Learning Check
What is the hybridization for the phosphorus atom in the molecule below?
Answer on the next slide
TO PREVIOUS SLIDE

44. UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
Learning Check
The electron-group arrangement for phosphorus is tetrahedral.
Therefore, P has an sp3 hybridization.
TO PREVIOUS SLIDE

45. The Influence of Molecular Shape on Polarity
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
The Influence of Molecular Shape on Polarity
The shape of a molecule affects that molecule’s polarity.
polar bonds have a bond dipole
bond dipoles are indicated using vectors that point in the direction of higher electron density
In a polar covalent bond, a partial positive charge is associated with one atom and a partial negative charge is associated with the other atom.
TO PREVIOUS SLIDE

46. Determining Whether a Molecule is Polar UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
Determining Whether a
Molecule is Polar
A molecule with one or more polar bonds is not necessarily a polar molecule. The molecule’s shape must be considered. The polarity as a whole can be determined by adding the vectors.
Both water and carbon dioxide have two polar bonds. But water’s bent shape results in a polar molecule, while carbon dioxide’s linear shape results in a non-polar molecule.
TO PREVIOUS SLIDE

47. Molecular Shapes and Polarities
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
Molecular Shapes and Polarities
TO PREVIOUS SLIDE

48. How Intermolecular Forces Affect the Properties of Solids and Liquids
UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
How Intermolecular Forces Affect the Properties of Solids and Liquids
Intermolecular forces exist between ions and molecules and influence the physical properties of substances.
Categories of forces:
dipole-dipole
ion-dipole
induced dipole
dispersion
TO PREVIOUS SLIDE

49. Dipole-Dipole UNIT 2 Dipole-dipole forces:
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
Dipole-Dipole
Dipole-dipole forces:
are forces of attraction between polar molecules, which have a region of partial positive charge and a region of partial negative charge
are a main reason for melting and boiling point differences between polar and non-polar molecules
include hydrogen bonding, as an example of one type
Hydrogen bonding
(dotted lines) in water
TO PREVIOUS SLIDE

50. Ion-Dipole UNIT 2 Ion-dipole forces:
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
Ion-Dipole
Ion-dipole forces:
are forces of attraction between partial charges on polar molecules and ions
depend on the size and charge of the ion and the magnitude of the partial charge and size of the molecule
are involved in the process of hydration
Ion-dipole intermolecular forces.
TO PREVIOUS SLIDE

51. Induced Dipoles UNIT 2 Dipole-induced dipole forces:
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
Induced Dipoles
Dipole-induced dipole forces:
are forces of attraction between a polar molecule and a non-polar molecule that has an induced (temporary) dipole due to the nearby polar molecule
Ion-induced dipole forces:
are forces of attraction between an ion and a non-polar molecule that has an induced dipole due to the nearby ion
A dipole can be induced in a non-polar molecule.
TO PREVIOUS SLIDE

52. Dispersion Forces UNIT 2 Dispersion forces:
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
Dispersion Forces
Dispersion forces:
are forces of attraction between all molecules, including non-polar molecules
are due to spontaneous temporary dipoles that form due to the constant motion of electrons in covalent bonds
depend on the size and shape of the molecules
the larger and more linear the molecule, the greater the force of attraction
The more linear molecule has a higher boiling point because the dispersion forces are greater.
TO PREVIOUS SLIDE

53. UNIT 2
Chapter 4: Chemical Bonding and Properties of Matter
Section 4.2
Section 4.2 Review
TO PREVIOUS SLIDE