Chapter 9 Chemical Bonds.

Slides:



Advertisements
Similar presentations
Chemical Bonding.
Advertisements

The Nature of Chemical Bonds
COVALENT BONDS Chapter 5 Section 3.
Formation of chemical bonds
Chemical bonds An introduction to chemistry. Compounds and chemical change Atom - smallest elemental unit Molecule –smallest particle still retaining.
An Introduction to Chemistry
Chemical Bonding Chemical bond – The attractive force between the protons of one atom for the electrons of another atom Determined by electronegativity.
Ionic and Covalent Bonds. Two or more elements chemically combined. Compound.
Chapter 9: Chemical Bonds. Chemical Terminology Atom Molecule Atom Molecule.
 Define these words  Ion  Ionic bond  Ionic compound  Chemical formula  Subscript  Covalent bond.
Pearson Prentice Hall Physical Science: Concepts in Action
Chapter 7 and 8.  Valence electrons are responsible for the bonding between two atoms.
Chemical Bonds Ionic and Covalent Bonding. Chemical Bonds – Ionic: Metals + Nonmetals – Covalent: Nonmetals + Nonmetals Sharing of electrons Mostly gases,
Unit 6A: Ionic and Covalent Bonding. Ions Why do elements in the same group behave similarly? They have the same number of valence electrons. Valence.
Pre AP Chemistry Chapter 6 “Chemical Bonding”. Introduction to Chemical Bonding Chemical bond – a mutual electrical attraction between the nuclei and.
Electron Dot Formulas Chemistry 7(C). Lesson Objectives Draw electron dot formulas – Ionic compounds – Covalent compounds Electron Dot Formulas.
Chapter #7 Chemical Bonds.. Chemical Bond An attractive force that holds two atoms together in a complex unit. Electrons combine to form chemical bonds.
CHAPTER 5: CHEMICAL BONDING Name:Prachayanee Chueamsuwanna Date: Oct. 19,2015.
Chapter 6: Chemical Bonds When the highest occupied energy level of an atom is filled with e, the atom is stable and not likely to react. In other words,
Chemical Bonding and Nomenclature
LECTURE 6: COVALENT COMPOUNDS.
Chapter 6 Chemical Bonds.
Ionic Compounds & Metals
Chapter 6 Objectives Section 1 Introduction to Chemical Bonding
CHEMICAL BONDING IONIC BONDS COVALENT BONDS HYDROGEN BONDS.
Unit 8 Bonding and Nomenclature
LESSON 2.2 Writing Formulas MgCl2.
Pearson Prentice Hall Physical Science: Concepts in Action
I. Introduction to Bonding
SECTION 1. INTRODUCTION TO CHEMICAL BONDING
Chapter 19 Chemical Bonds.
Chemical Bonding.
Remember - Ions are atoms or groups of atoms with a charge
Unit 10 – Chemical Bonding & Nomenclature
Chemical Bonds.
Chapter 6 – Chemical Bonds
Chapter 7: Covalent Bonding and Lewis Formula’s
Ionic Bonds.
Intramolecular Forces Intermolecular Forces
Formation of Ionic Compounds
The Structure of matter
CHEMICAL BONDS.
How Elements Form Compounds
Chapter 6 – Chemical Bonds
Chemical Bonding.
Chapter 8 Basic Concepts of Chemical Bonding
Chemical Bonds.
Intramolecular Forces Intermolecular Forces
Chapter 6 Table of Contents Chemical Bonding
Chemical Bonding Unit 2 Topic 3 Chapter 6.
Chemical Bonding.
Chemical Bonding.
Bonding.
Chapter 6 Objectives Define chemical bond.
Chapter 6 Chemical Bonds
Chemical bonding Chapter 22 Section 2 Pages
6.4 LEWIS STRUCTURE DIAGRAMS
Forming Chemical Bonds
Chemical Bonding.
Chapter 6: Chemical Bonds
Chapter 6- Chemical Bonding
Pearson Prentice Hall Physical Science: Concepts in Action
Chemical Bonds.
Chemical Bonding.
Electron Configurations – a Review and More…
Chemistry Mrs. Partridge
Bonding – Introduction May 12
Names and Formulas for Acids
Chapters 7 and 8 – Bonding.
…electrons are transferred
Presentation transcript:

Chapter 9 Chemical Bonds

Core Concept Electron structure will explain how and why atoms join together in certain numbers.

Valence Electrons and Ions Outer electrons determine the chemical properties of an atom Octet rule Atoms attempt to acquire an outer shell of eight electrons Electrons can be gained/lost/shared in the process Example: sodium (Na)

Chemical Bonds Covalent Three types: Ionic Metallic bonds Attractive forces holding atoms together in compounds Three types: Ionic Electrons transferred between atoms Electrostatic force = binding force Covalent Octets achieved through sharing electrons Typically between nonmetallic elements Metallic bonds Outer electrons move freely throughout metal “Electron sea” within rigid lattice of metal atoms Conduct heat and electricity well

Ionic Bonds Chemical bond of electrostatic attraction Forms between metals and nonmetals Form crystalline solids with orderly geometric structure Example: NaCl Na loses; Cl gains No single NaCl molecule, per se

Ionic Compound Formulas Ionic compounds have an overall neutral charge If charges aren’t complementary on ions, multiple ions are needed Multiple ions are represented by subscripts

Ionic Compound Formulas Two rules Write symbol for positive ion first followed by negative ion symbol Assign subscripts to assure compound is electrically neutral Example: Calcium chloride

Ionic Compound Names Name of metal (positive) ion first Name is unchanged Stem name of second element next ending in “-ide” (for two elements) Many elements have variable charges Historical suffix usage “-ic” for higher of two; “-ous” for lower Modern approach English name of metal followed by Roman numeral indicating charge

Lewis Dot Structures for Ionic Compounds For anions fill the octet by adding the required number of electrons Example O atom becoming O2-. For cations empty the octet by losing all electrons from the valence shell. Example Mg atom becoming Mg2+.

Lewis Dot Structures for Ionic Compounds Place squared brackets around each ion and place the charge as a superscript following the brackets. To represent the number of each ion that is needed place a coefficient in front of the squared brackets for each ion to represent the required quantity. Ex: Calcium chloride [Ca]+2 + 2[Cl]-1

Covalent Bonds Chemical bonds formed by sharing pairs of electrons Formed between 2 nonmetals Electrons shared to form octets, ideally Overlap of shared electron clouds between nuclei yields net attraction Atoms within covalent compounds are electrically neutral, or nearly so

Covalent Compounds and Formulas Covalent compound - held together by covalent bonds Electrons shared in covalent bonds Electron dot representation Bonding pairs shared Lone (non-bonding) pairs not shared

Multiple Bonds Sharing of more than one electron pair Examples Ethylene - double bond Acetylene - triple bond

Composition of Compounds Millions of different combinations of over 90 elements Common names Often related to historical usage (baking soda, washing soda,…) Difficult to relate to actual molecular composition Modern approach - systematic sets of rules Different for ionic and covalent compounds One common rule - “-ide” means compound contains only two different elements

Covalent Compound Names Molecular - composed of two or more nonmetals Same elements can combine to form a number of different compounds Two rules First element in formula named first with number indicated by Greek prefix Stem name of second element next; Greek prefix for number; ending in “-ide” (for two elements)

Lewis Structures Lewis structures are representations of molecules showing all valence electrons, bonding and nonbonding.

Electronegativities

Writing Lewis Structures Find the sum of valence electrons of all atoms in the polyatomic ion or molecule. If it is an anion, add one electron for each negative charge. If it is a cation, subtract one electron for each positive charge. PCl3 5 + 3(7) = 26

Writing Lewis Structures The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds. Keep track of the electrons: 26  6 = 20

Writing Lewis Structures Fill the octets of the outer atoms. Keep track of the electrons: 26  6 = 20  18 = 2

Writing Lewis Structures Fill the octet of the central atom. Keep track of the electrons: 26  6 = 20  18 = 2  2 = 0

Writing Lewis Structures If you run out of electrons before the central atom has an octet… …form multiple bonds until it does.

Bond Polarity Result of unequal sharing of electrons Electronegativity Measure of an atom’s ability to attract electrons Differences: 1.7 or greater - ionic 0.5-1.7 - polar covalent Less than 0.5 - covalent