Chapter 5 Compounds and Their Bonds

Slides:



Advertisements
Similar presentations
Unit 4 Test Review.
Advertisements

Ions and Ionic Bonds.
1 4.6 Covalent Compounds Copyright © 2005 by Pearson Education, Inc. Publishing as Benjamin Cummings Chapter 4 Forces Between Particles.
1 4.1 Valence Electrons 4.2 Octet Rule and Ions Chapter 4 Compounds and Their Bonds.
BONDING Ch 8 & 9 – Honors Chemistry General Rule of Thumb:
More bonding Quick Overview of: Ionic Bonding Metallic bonding Hydrogen bonding Quick Overview of: Ionic Bonding Metallic bonding Hydrogen bonding.
Chapter 4 Octet Rule and Ions
Copyright © Houghton Mifflin Company. All rights reserved. 12 | 1 Chemical Bonds Forces that hold atoms together Ionic bonds: the forces of attraction.
Chapter 6 Ionic Compounds
Chapter 7 Molecular Structure: Solids and Liquids Electron Configuration of Ionic Compounds Review.
Compounds and Their Bonds Covalent Bonds Covalent Compounds Bond Polarity Polyatomic Ions.
Chapter 6 Ionic Compounds Copyright © 2008 by Pearson Education, Inc. Publishing as Benjamin Cummings Have your i-clickers ready Silence cell phones and.
More bonding Quick Overview of: Ionic Bonding Metallic bonding
Unit 04 Chemical Bonding.
Chapter 4 Compounds and Their Bonds
Compounds and their Bonds
Chapter 4 Compounds and Their Bonds
Forces that hold atoms together.  There are several major types of bonds. Ionic, covalent and metallic bonds are the three most common types of bonds.
Copyright © 2004 Pearson Education Inc., publishing as Benjamin Cummings Covalent Bonds 4.6 Naming and Writing Formulas of Covalent Compounds 4.7.
What are compounds and how they form bonds? Octet Rule and Ions An octet is 8 valence electrons. is associated with stability of noble gases. He is stable.
Chapter 4 Compounds and Their Bonds
Chapter 7 Molecular Structure: Solids and Liquids
Chapter 6 Molecules and Covalent Compounds Copyright © 2005 by Pearson Education, Inc. Publishing as Benjamin Cummings.
Chapter 5 Compounds and Their Bonds
BONDING Chapters 4 & 12.
Chapter 6 Ionic Compounds Copyright © 2008 by Pearson Education, Inc. Publishing as Benjamin Cummings Have your i-clickers ready Silence cell phones and.
Chapter 4 Compounds and Their Bonds 4.4 Polyatomic Ions 1.
Ch. 3 HW- 3.18, 3.21, 3.32, 3.33, 3.38, 3.39, 3.43, 3.52, 3.53, 3.56, 3.59, 3.61.
Chemistry B2A Chapter 12 Chemical Bonding.
Chemical Formulas and Names for Ionic Compounds
Bettelheim, Brown, Campbell and Farrell Chapter 3
Metals  Lose e-s  positive ion
Chapter 6 Ionic Compounds Copyright © 2008 by Pearson Education, Inc. Publishing as Benjamin Cummings Silence cell phones and pagers.
Chapter 5 Ionic Compounds Copyright © 2005 by Pearson Education, Inc. Publishing as Benjamin Cummings.
Chapter 4 Compounds and Their Bonds 4.1 Octet Rule and Ions 1 Copyright © 2009 by Pearson Education, Inc.
1 Chapter 4 Compounds and Their Bonds 4.1 Octet Rule and Ions Copyright © 2005 by Pearson Education, Inc. Publishing as Benjamin Cummings.
Chapter 4 Compounds and Their Bonds 4.1 Octet Rule and Ions 1.
Chemical Bonds I. Why Atoms Combine  Chemical Formula  Chemical Bond  Stability.
Chemical Bonds I. Why Atoms Combine  Chemical Formulas  Chemical Bonds  Stability.
Chapter 4 Covalent Compounds.
Copyright © 2004 Pearson Education Inc., publishing as Benjamin Cummings Valence Electrons 4.2 Octet Rule and Ions Chapter 4 Compounds and Their.
Chemical Bonds I. Why Atoms Combine  Chemical Formulas  Chemical Bonds  Stability.
1 4.4 Polyatomic Ions Chapter 4 Compounds and Their Bonds Copyright © 2009 by Pearson Education, Inc.
Chapter 4 Compounds and Their Bonds 4.5 Covalent Compounds 1 Copyright © 2009 by Pearson Education, Inc.
Chapter 10 Molecular Structure: Solids and Liquids Electron Configurations and Dot Formulas Review.
Reactivity of Elements How many valence electrons do the noble gases have? 8 electrons = octet Compounds form when two or more elements come together,
Chapter 4 Compounds and Their Bonds 4.4 Polyatomic Ions 1 Copyright © 2009 by Pearson Education, Inc.
Chemical Bonds. Chapter 52 Ions Ions have different numbers of electrons and protons Cations have lost electrons and have net positive charge Anions have.
Section 12.1 Characteristics of Chemical Bonds 1.To learn about ionic and covalent bonds and explain how they are formed 2.To learn about the polar covalent.
Covalent Compounds Chapter 8. Section 1, Covalent Bonds –Remember, ionic compounds are formed by gaining and losing electrons –Atoms can also share electrons.
Chemical Bonds I. Why Atoms Combine  Chemical Formulas  Chemical Bonds  Stability.
DE Chemistry – King William High School.  Cation – positive charge  lose an electron  Alkali metals (group 1) take a +1 charge  Alkaline earth metals.
1 Chemical Bonds The Formation of Compounds From Atoms Chapter 11 Hein and Arena.
a. protons b. neutrons c. electrons d. morons a. protons b. neutrons c. electrons d. morons.
Unit 5: Chemical Bonding Chapters 8 & 9 Test - November 21, 2008.
Chapter 3 Lecture General, Organic, and Biological Chemistry: An Integrated Approach Laura Frost, Todd Deal and Karen Timberlake by Richard Triplett Chapter.
Basic Chemistry Copyright © 2011 Pearson Education, Inc. 1 Chapter 6 Inorganic and Organic Compounds: Names and Formulas 6.1 Octet Rule and Ions Basic.
SOL Review 3 Bonding and Naming Ionic and Covalent Compounds.
The 8 valance electrons in the noble gases make them chemically stable All other Elements “want” their valence electron structure to look like a noble.
Basic Chemistry Copyright © 2011 Pearson Education, Inc. Chemistry Chapter 6 Key Vocabulary: 1.octet 2.Octet rule 3.Duet 4.Ionic bond 5.Covalent bond 6.cation.
Unit 6: Chemical Bonding and Intermolecular Forces
Chapter 4 Compounds and Their Bonds
Chapter 4 Compounds and Their Bonds
Chapter 6 Ionic Compounds
H2O A. Chemical Formula Shows: 1) elements in the compound
1.3 Ions and Octet Rule.
Chapter 4 Compounds and Their Bonds
Chapter 7 Molecular Structure: Solids and Liquids
CHEMICAL BONDING Cocaine
Ch Chemical Bonds I. Why Atoms Combine (p ) Chemical Formula
Presentation transcript:

Chapter 5 Compounds and Their Bonds 5.1 Octet Rule and Ions

Octet Rule An octet is 8 valence electrons is associated with the stability of the noble gases does not occur with He; He is stable with 2 valence electrons (duet) Valence Electrons He 1s2 2 Ne 1s22s22p6 8 Ar 1s22s22p63s23p6 8 Kr 1s22s22p63s23p64s23d104p6 8

Ionic and Covalent Bonds Atoms form octets to become more stable by losing, gaining, or sharing valence electrons by forming ionic or covalent bonds

Metals Form Positive Ions by a loss of their valence electrons with the electron configuration of the nearest noble gas that have fewer electrons than protons Group 1A(1) metals ion 1+ Group 2A(2) metals ion 2+ Group 3A(3) metals ion 3+

Formation of a Sodium Ion, Na+ Sodium achieves an octet by losing its one valence electron. With the loss of its valence electron, a sodium ion has a 1+ charge. Sodium atom Sodium ion 11p+ 11p+ 11e– 10e– 0 1 +

Formation of Magnesium Ion, Mg2+ Magnesium achieves an octet by losing its two valence electrons. With the loss of two valence electrons magnesium forms a positive ion with a 2+ charge Mg atom Mg2+ ion 12p+ 12p+ 12e– 10e– 0 2+

Formation of Negative Ions In ionic compounds, nonmetals achieve an octet arrangement gain electrons form negatively charged ions with 3–, 2–, or 1– charges

Formation of a Chloride Ion, Cl– Chlorine achieves an octet by adding an electron to its valence electrons.

Ionic Charge from Group Numbers The charge of a positive ion is equal to its Group number. Group 1A(1) = 1+ Group 2A(2) = 2+ Group 3A(3) = 3+ The charge of a negative ion is obtained by subtracting 8 or 18 from its Group number . Group 6A(16) = 6 – 8 = 2– or 16 – 18 = 2–

Group Number and Ionic Charge Ions achieve the electron configuration of their nearest noble gas of metals in Groups 1A(1), 2A(2), or 3A(13) have positive 1+, 2+, or 3+ charge. Of nonmetals in Groups 5A(15), 6A(16), or 7A(17) have negative 3–, 2–, or 1– charge.

Groups Numbers for Some Positive and Negative Ions

Chapter 5 Compounds and Their Bonds 5.2 Ionic Compounds

Ionic Compounds Ionic compounds consist of positive and negative ions have attractions called ionic bonds between positively and negatively charged ions have high melting and boiling points are solid at room temperature

Salt Is an Ionic Compound Sodium chloride, or “table salt,” is an example of an ionic compound.

Ionic Formulas An ionic formula consists of positively and negatively charged ions is neutral has charge balance total positive charge = total negative charge The symbol of the metal is written first, followed by the symbol of the nonmetal.

Charge Balance for NaCl, “Salt” In NaCl, a Na atom loses its valence electron a Cl atom gains an electron the symbol of the metal is written first, followed by the symbol of the nonmetal.

Charge Balance In MgCl2 In MgCl2, a Mg atom loses two valence electrons two Cl atoms each gain one electron subscripts indicate the number of ions needed to give charge balance

Writing Ionic Formulas from Charges Charge balance is used to write the formula for sodium nitride, a compound containing Na+ and N3−. Na+ 3 Na+ + N3− = Na3N 3(1+) + 1(3–) = 0

Chapter 5 Compounds and Their Bonds 5.3 Naming and Writing Ionic Formulas

Naming Ionic Compounds with Two Elements To name a compound with two elements, identify the cation and anion name the cation first, followed by the name of the anion

Examples of Ionic Compounds with Two Elements Formula Ions Name Cation Anion NaCl Na+ Cl– sodium chloride K2S K+ S2– potassium sulfide MgO Mg2+ O2– magnesium oxide CaI2 Ca2+ I– calcium iodide Al2O3 Al3+ S2– aluminum sulfide

Transition Metals Form Positive Ions Most transition metals and Group 4(14) metals, Form 2 or more positive ions Zn2+, Ag+, and Cd2+ form only one ion.

Metals with Variable Charge The names of transition metals with two or more positive ions (cations) use a Roman numeral after the name of the metal to identify ionic charge.

Naming Ionic Compounds with Variable Charge Metals

Naming FeCl2 STEP 1 Determine the charge of the cation from the anion. Fe ion + 2 Cl– = Fe ion + 2– = 0 Fe ion = 2+ = Fe2+ STEP 2 Name the cation by the element name, and use a Roman numeral to show its charge. Fe2+ = iron(II) STEP 3 Write the anion with an ide ending. chloride STEP 4 Name the cation first, then the anion. iron(II) chloride

Naming Cr2O3 STEP 1 Determine the charge of cation from the anion. 2Cr ions + 3O2– = 2Cr ions + 3(2–) = 2Cr ions + 6– = 0 2Cr ions = 6+ Cr ion = 3+ = Cr3+ STEP 2 Name the cation by the element name, and use a Roman numeral to show its charge. Cr3+ = chromium(III) STEP 3 Write the anion with an ide ending. oxide STEP 4 Name the cation first, then the anion. chromium (III) oxide

Guide to Writing Formulas from the Name

Writing Formulas Write a formula for potassium sulfide. STEP 1 Identify the cation and anion. potassium = K+ sulfide = S2− STEP 2 Balance the charges. K+ S2− K+ 2(1+) + 1(2–) = 0 STEP 3 Write the cation first. 2K+ and 1S2− = K2S1 = K2S

Writing Formulas STEP 1 Identify the cation and anion. Write a formula for iron(III) chloride. STEP 1 Identify the cation and anion. iron (III) = Fe3+ (III = charge of 3+) chloride = Cl− STEP 2 Balance the charges. Fe3+ Cl− Cl− 1(3+) + 3(1–) = 0 STEP 3 Write the cation first. 1Fe3+ and 3Cl− = FeCl3

Chapter 5 Compounds and Their Bonds 5.4 Polyatomic Ions

Polyatomic Ions A polyatomic ion is a group of atoms has an overall ionic charge Examples: NH4+ ammonium OH− hydroxide NO3− nitrate NO2− nitrite CO32− carbonate PO43− phosphate HCO3− hydrogen carbonate (bicarbonate)

Some Names of Polyatomic Ions The names of common polyatomic anions end in ate NO3− nitrate PO43− phosphate with one oxygen less end in ite NO2− nitrite PO33− phosphite with hydrogen attached use prefix hydrogen (or bi) HCO3− hydrogen carbonate (bicarbonate) HSO3− hydrogen sulfite (bisulfite)

Guide to Naming Compounds with Polyatomic Ions

Examples of Names of Compounds with Polyatomic Ions The positive ion is named first followed by the name of the polyatomic ion. NaNO3 sodium nitrate K2SO4 potassium sulfate Fe(HCO3)3 iron(III) bicarbonate or iron(III) hydrogen carbonate (NH4)3PO3 ammonium phosphite

Writing Formulas with Polyatomic Ions The formula of an ionic compound containing a polyatomic ion must have a charge balance that equals zero (0) Na+ and NO3− NaNO3 with two or more polyatomic ions put the polyatomic ions in parentheses. Mg2+ and 2NO3− Mg(NO3)2 subscript 2 for charge balance

Chapter 5 Compounds and Their Bonds 5.5 Covalent Compounds

Forming Octets in Molecules In a fluorine (F2) molecule, each F atom shares one electron acquires an octet

Diatomic Elements These elements share electrons to form diatomic, covalent molecules.

Guide to Writing Electron-Dot Formulas

Single and Multiple Bonds In a single bond, one pair of electrons is shared. In a double bond, two pairs of electrons are shared. In a triple bond, three pairs of electrons are shared.

Electron-Dot Formula of CS2 Write the electron-dot formula for CS2. STEP 1 Determine the atom arrangement. The C atom is the central atom. S C S STEP 2 Determine the total number of valence electrons for 1C and 2S. 1 C(4e–) + 2 S(6e–) = 16e– STEP 3 Attach each S atom to the central C atom using one electron pair. S : C : S 16e– – 4e– = 12e– remaining STEP 4 Attach 12 electrons as 6 lone pairs. .. .. : S : C : S :

Electron-Dot Formula of CS2 (continued) To complete octets, form one or more multiple bonds. Convert two lone pairs to bonding pairs between C and S atoms to make two double bonds.

A Nitrogen Molecule has A Triple Bond In a nitrogen molecule, N2, each N atom shares 3 electrons each N atom attains an octet the sharing of 3 sets of electrons is a multiple bond called a triple bond

Resonance Structures Resonance structures are two or more electron-dot formulas for the same arrangement of atoms related by a double-headed arrow ( ) written by changing the location of a double bond between the central atom and a different attached atom

Writing Resonance Structures Sulfur dioxide has two resonance structures. STEP 1 Write the arrangement of atoms. O S O STEP 2 Determine the total number of valence electrons. 1 S(6e−) + 2 O(6e−) = 18e− STEP 3 Connect bonded atoms by single electron pairs. O : S : O 4e− used 18e− – 4e− = 14e− remaining

Writing Resonance Structures (continued) STEP 4 Add 14 remaining electrons as 7 lone pairs. STEP 5 Form a double bond to complete octets. Two resonance structures are possible.

Formula from Ionic Charges Write the ionic formula of the compound containing Ba2+ and Cl. Write the symbols of the ions. Ba2+ Cl Balance the charges. Ba2+ Cl two Cl needed Cl Write the ionic formula using a subscript 2 for two chloride ions that give charge balance. BaCl2

Chapter 5 Compounds and Their Bonds 5.6 Naming and Writing Covalent Formulas NO nitrogen oxide NO2 nitrogen dioxide N2O4 dinitrogen tetroxide

Names of Covalent Compounds Prefixes are used in the names of covalent compounds because two nonmetals can form two or more different compounds Examples of compounds of N and O: NO nitrogen oxide NO2 nitrogen dioxide N2O dinitrogen oxide N2O4 dinitrogen tetroxide N2O5 dinitrogen pentoxide

Naming Covalent Compounds STEP 1 Name the first nonmetal by its element name. STEP 2 Name the second nonmetal with an ide ending. STEP 3 prefixes to indicate the number (from subscripts) of atoms of each nonmetal. Mono is usually omitted.

Naming Covalent Compounds (Continued) What is the name of SO3? STEP 1 The first nonmetal is S sulfur. STEP 2 The second nonmetal is O, named oxide. STEP 3 The subscript 3 of O is shown as the prefix tri. SO3 → sulfur trioxide The subscript 1(for S) or mono is understood.

Naming Covalent Compounds (Continued) Name P4S3 STEP 1 The first nonmetal P is phosphorus. STEP 2 The second nonmetal S is sulfide. STEP 3 The subscript 4 of P is shown as tetra. The subscript 3 of O is shown as tri. P4S3 → tetraphosphorus trisulfide

Guide to Writing Formulas for Covalent Compounds

Writing Formulas of Covalent Compounds Write the formula for carbon disulfide. STEP 1 Elements are C and S STEP 2 No prefix for carbon means 1 C Prefix di = 2 Formula: CS2

Chapter 5 Compounds and Their Bonds 5.7 Electronegativity and Bond Polarity

Electronegativity The electronegativity value indicates the attraction of an atom for shared electrons increases from left to right going across a period on the periodic table decreases going down a group on the periodic table is high for the nonmetals, with fluorine as the highest is low for the metals

Some Electronegativity Values High values Low values

Nonpolar Covalent Bonds A nonpolar covalent bond occurs between nonmetals has an equal or almost equal sharing of electrons has almost no electronegativity difference (0.0 to 0.4) Examples: Atoms Electronegativity Type of Bond Difference _______________ N–N 3.0 – 3.0 = 0.0 Nonpolar covalent Cl–Br 3.0 – 2.8 = 0.2 Nonpolar covalent H–Si 2.1 – 1.8 = 0.3 Nonpolar covalent

Polar Covalent Bonds A polar covalent bond occurs between nonmetal atoms has an unequal sharing of electrons has a moderate electronegativity difference (0.5 to 1.7) Examples: Atoms Electronegativity Type of Bond Difference _________________________________________ O–Cl 3.5 – 3.0 = 0.5 Polar covalent Cl–C 3.0 – 2.5 = 0.5 Polar covalent O–S 3.5 – 2.5 = 1.0 Polar covalent

Comparing Nonpolar and Polar Covalent Bonds

Chapter 5 Compounds and Their Bonds 5.8 Shapes and Polarity of Molecules

VSEPR In the valence-shell electron-pair repulsion theory (VSEPR), the electron groups around a central atom are arranged as far apart from each other as possible have the least amount of repulsion of the negatively charged electrons have a geometry around the central atom that determines molecular shape

Shapes of Molecules The three-dimensional shape of a molecule is the result of bonded groups and lone pairs of electrons around the central atom is predicted using the VSEPR theory (valence-shell-electron-pair repulsion)

Guide to Predicting Molecular Shape (VSEPR Theory)

Two Electron Groups In a molecule of BeCl2, there are two electron groups bonded to the central atom, Be (Be is an exception to the octet rule) .. .. : Cl : Be : Cl : to minimize repulsion, the arrangement of two electron groups is 180°, or opposite each other the shape of the molecule is linear

Two Electron Groups with Double Bonds In a molecule of CO2, there are two electron groups bonded to C (electrons in each double bond are counted as one group) repulsion is minimized with the double bonds opposite each other at 180° the shape of the molecule is linear

Three Electron Groups In a molecule of BF3, three electron groups are bonded to the central atom B (B is an exception to the octet rule) .. : F: .. .. .. : F : B : F : .. .. repulsion is minimized with 3 electron groups at angles of 120° the shape is trigonal planar

Two Electron Groups and One Lone Pair In a molecule of SO2, S has 3 electron groups; 2 electron groups bonded to O atoms and one lone pair .. .. .. :O :: S : O : .. repulsion is minimized with the electron groups at angles of 120°, a trigonal planar arrangement the shape is bent (120°), with two O atoms bonded to S ● ●

Four Electron Groups In a molecule of CH4, there are four electron groups around C repulsion is minimized by placing four electron groups at angles of 109°, which is a tetrahedral arrangement the four bonded atoms form a tetrahedral shape

Three Bonding Atoms and One Lone Pair In a molecule of NH3, three electron groups bond to H atoms, and the fourth one is a lone (nonbonding) pair repulsion is minimized with 4 electron groups in a tetrahedral arrangement the three bonded atoms form a pyramidal (~109°) shape

Two Bonding Atoms and Two Lone Pairs In a molecule of H2O, two electron groups are bonded to H atoms and two are lone pairs (4 electron groups) four electron groups minimize repulsion in a tetrahedral arrangement the shape with two bonded atoms is bent (~109°)

Determining Molecular Polarity Determine the polarity of the H2O molecule. Solution: The four electron groups of oxygen are bonded to two H atoms. Thus the H2O molecule has a net dipole, which makes it a polar molecule.

Chapter 5 Compounds and Their Bonds 5.9 Attractive Forces in Compounds

Dipole–Dipole Attractions In covalent compounds, polar molecules exert attractive forces called dipole-dipole attractions form strong dipole attractions called hydrogen bonds between hydrogen atoms bonded to F, O, or N, and other atoms that are very electronegative

Dipole–Dipole Attractions (continued)

Dispersion Forces Dispersion forces are weak attractions between nonpolar molecules caused by temporary dipoles that develop when electrons are not distributed equally

Comparison of Bonding and Attractive Forces

Melting Points and Attractive Forces Ionic compounds require large amounts of energy to break apart ionic bonds. Thus, they have high melting points. Hydrogen bonds are the strongest type of dipole–dipole attractions. They require more energy to break than do other dipole attractions. Dispersion forces are weak interactions, and very little energy is needed to change state.