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

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

4.1 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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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