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  2. Purpose • To be able to understand the basic concept of bonding in metal atoms, nonmetal atoms, and ions. • Learn how to draw accurate Lewis structure and indicate its formal charge. • Know how to name the molecules. • Know how to calculate the bond enthalpy.

  3. Bonding • A chemical bond is an attraction between atoms that allows the formation of chemical substances that contain two or more atoms. • Three basic types of bonds includes • Ionic- Electrostatic attraction between ions. • Covalent- Sharing of electrons. • Metallic- Metal atoms bonded to several other atoms.

  4. Valence Electrons • Valence electrons are electrons in the outermost shell of an atom; electrons in the highest occupied energy level. • The number of valence electrons determines the properties of an element.

  5. Lewis Dot Symbols • G.N. Lewis postulated that atoms achieve greater stability by bonding and procuring a noble gas electronic structure. • When forming compounds, atoms tend to add or subtract electrons until they are surrounded by eight valence electrons (the octet rule). • He devised a system, the Lewis dot system, where dots represent the valence electrons, and are moved or combined to represent processes and bonding.

  6. Examples

  7. Formation of Cations • A positively charged ion called a cation is produced when an atoms loses one or more valence electrons. • e.g. Li → Li⁺ + e¯

  8. Formation of Anion • A negatively charged ion is called anion and is produced when an atoms loses electrons. Note that the name of anions ends with –ide. • Cl + e¯ → Cl¯ chloride

  9. Formation of Ionic Bonds & Ionic Compounds • An ionic compound is one formed between a cation and an anion. Although they are composed of ions, ionic compound are electrically neutral. • Cation and anions have opposite charges and therefore attracts each other by means of electrostatic forces forming bonds called ionic bonds. • Because compounds are electrically neutral.

  10. Naming Cation • Cations (except ammonium NH4+) derive their names unchanged from the metal they are derived from. When cations occur with two different charges the suffix –ous and –ic were used to designate the lower and higher oxidation states respectively.

  11. Naming Anion • Anions take the elements name but with the suffix changed to –ide.

  12. Naming Oxyanion Compounds • The one with the fewest oxygens has the prefix hypo- and ends in –ite (ClO−: hypochlorite). • The one with the second fewest oxygens ends in –ite (ClO₂‾: chlorite). • The one with the second most oxygens ends in –ate (ClO₃‾: chlorate). • The one with the most oxygens has the prefix per- and ends in –ate (ClO₄‾: perchlorate).

  13. Lattice Energy • Lattice Energy is the energy required to completely separate a mole of a solid ionic compound into its gaseous ions. • Lattice energy, then, increases with the charge on the ions. • It also increases with decreasing size of ions.

  14. Lattice Energy Table

  15. Properties of Ionic Compounds • Exist in solid state. • Form crystal lattices not molecules. • Good insulators. • High melting points/ Boiling Points. • Conduct electricity when dissolved in water or as a liquid; are strong electrolytes. Solids do not conduct electricity. • Are insoluble in the organic compounds. • Are very reactive. • Most are brittle and break under stress.

  16. Metallic Bonding • Metals consists of closely packed cations and losely held valence elctrons. • The valence electrons in the metal can be modeled as a sea of electrons. The valence electrons are mobile and can drift from one part of the metal to the other. • Metallic bond are forces of attraction between the free-floating delocalized valence electrons and the positively charged metal ions. The bonds hold metals together.

  17. Properties of Metals • Conduct heat & conduct electricity • Generally high melting and boiling points • Strong • Malleable (can be hammered or pressed out of shape without breaking) • Ductile (able to be drawn into a wire) • Metallic luster • Opaque (reflect light)

  18. Alloys • An alloy is a mixture or metallic solid solution composed of two or more elements. • Alloys are important because their properties are often superior to those of their components.

  19. Examples • Duralumin - copper and aluminium • Woods metal - lead, tin and cadmium • Bronze - copper and tin • Brass - copper and zinc • Rose gold - copper and gold • Solder - lead and tin • Steel - iron and carbon, often other metals as well such as Mn, B, and Cr. • Nichrome(chromium, iron, nickel) • Surgical stainless steel ( iron, carbon, chromium, molybdenum, nickel)

  20. Covalent Bond

  21. Molecules & Molecular Compounds • Ionic bonds are formed by giving up and accepting electrons. Atoms that are held together by sharing electrons are joined by a covalent bond. • A molecule is comprised of atoms held together by chemical bonds. • Molecular compounds are composed of different atoms bonded together by covalent bonds and almost always contain only nonmetals. • Polyatomic molecules are molecules where more than two atoms are bonded.

  22. Representing Molecules • A molecular formula shows how many atoms of each element a substance contain • Structural formulas show the order in which atoms are bonded. • Perspective drawings also show the three-dimensional array of atoms in a compound. • We use molecular models to try and visualise and represent molecules. These include space filling molecular model and ball-and-stick molecular model.

  23. Lewis Structures • Lewis suggested that a bond between atoms is the sharing of electrons, called a covalent bond, represented in Lewis structures as: • Molecules that are comprised of covalent bonds only are called covalent compounds. When there are many valence electrons, only the unpaired electrons are involved in bonding, the others remain unbonded and are called lone pairs or unshared pair.

  24. Covalent Bonds • Procuring 8 electrons is called the octet rule. • Single bonds are where two atoms share one pair of electrons. • Double bonds when two atoms share two pairs of electrons. • Triple bonds when three pairs are shared. • Multiple bonds are more stable than single bonds. • There are several electrostatic interactions in these bonds • Attractions between electrons and nuclei, • Repulsions between electrons, • Repulsions between nuclei.

  25. Coordinate (Dative) Covalent Bonds • In a coordinate covalent bond, the shared electron pair comes from one of the bonding atoms.

  26. Naming Molecular Compounds • Nomeclature is similar to ionic compounds. Greek multiplicity prefixes used to indicate multiplicity. The mono-prefix is not used for the first atom, and in oxides of the multiplicity prefix is left out. With hydrides, trivial names, and non-prefixed names are in common use. • Naming • Binary Compounds • Acids and Bases • Hydrates • Organic Compounds

  27. Naming Binary Compounds Nomenclature • The ending on the more electronegative element is changed to –ide. (CO₂: carbon dioxide and CCl₄: carbon tetrachloride). • If the prefix ends with a or o and the name of the element begin with a vowel, the two successive vowels are often elided into one (N₂O₅: dinitrogenpentoxide).

  28. Naming Acids and Bases Nomenclature • An acid produces a hydrogen cation (proton) when dissolved n water. • Anions whose names end in –ide form acids that are prefixed with “hydro” and the anion’s suffix becomes “ic.” • Oxoacidsare those that have a central atom, oxygens, and hydrogens toward the outside. The main acid usually has the “ic” suffix. Adding an O changes the name to “per…ic” acid. • Removing an O gives “-ous” acid (e.gHClO₂: chlorous acid), and removing two O’s produces “hypo…ous” acid. • Rules for naming anions: • If all the protons are removed from the normal “-ic” acid, the –“-ate” suffix serves. If all the protons from the “–ous” acid is removed the suffix is “-ite.”

  29. Naming Hydrates • Hydrates are compounds that bind water in when crystallised. • They are named with the number of water molecules indicated afterwards. • The hydrates is name by naming the cation first, followed by “hydrates” with prefix to indicate the number of water molecules. • E.g.CuSO45H2O is named copper (II) sulphatepentahydrate

  30. Naming Organic Compounds Nomenclature • Organic chemistry is the study of carbon. • Straight chain hydrocarbons are named with the Greek numerical prefix for the number of carbon atoms, and the suffix “-ane.” • The first four, do not use the Greek prefix: Methane, ethane, propane and butane.

  31. Bond Polarity • In a covalent bond the electrons are shared between two equal atoms. • In an ionic bond the electrons are practically transferred completely from one atom to the other, creating ions of opposite polarity. • Bond polarity is a measure of how equally or unequally the electrons in any covalent bond are shared. • A nonpolar covalent bond is one in which the electrons are shared equally, as is Cl₂ and N₂. • A polar covalent bond is one in which one of the atoms exerts a greater attraction for the bonding electrons than the other.

  32. Electronegativity • Atoms of widely varying electronegativities form ionic bonds, and those of closer electronegativities form polar covalent bonds. • Atoms in a bond with an electonegativity difference greater than 2.0 are ionic and less, polar covalent.

  33. Dipole Moments • Dipole Moments – Occurs when electrons are not shared equally in a covalent bond. When two atoms share electrons unequally, a bond dipole results. • The dipole moment, μ, produced by two equal but opposite charges separated by a distance, r, is calculated: μ = Qr • It is measured in debyes (D). • This is indicated by a crossed arrow, and leads to partial charges at the various atoms designated by a arrow to indicate the charge is partial:

  34. Attraction between Molecules • Intermolecular attractions are weaker than either ionic or covalent bonds. These include • Van der Waals Forces - the attraction of intermolecular forces between molecules.  • Dipole Interactions - two dipolar molecules interact with each other through space. • Dispersion forces - Interactions between ions, dipoles, and induced dipoles account for many properties of molecules  • Hydrogen Bonds - special type of dipole-dipole attraction which occurs when a hydrogen atom bonded to a strongly electronegative atom exists in the vicinity of another electronegative atom with a lone pair of electrons.

  35. Drawing the Lewis Structure • 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. • e.g. in PCl₃, # of valence electrons is: 5 + 3(7) = 26 • 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 • Fill the octets of the outer atoms. Keep track of the electrons: 26 − 6 = 20; 20 − 18 = 2 • Fill the octet of the central atom. Keep track of the electrons: 26 − 6 = 20; 20 − 18 = 2; 2 − 2 = 0 • If you run out of electrons before the central atom has an octet…form multiple bonds until it does.

  36. Formal Charges • For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms. • Subtract that from the number of valence electrons for that atom: the difference is its formal charge.

  37. Resonance Structure • Resonance means when two or more Lewis structures are required to represent a compound. Resonance structures are imaginary models to try and describe the real thing which is somewhere between.

  38. Exceptions to the Octet Rule • Ions or molecules with an odd number of electrons. • Ions or molecules with less than an octet. • Ions or molecules with more than eight valence electrons (an expanded octet).

  39. Bond Energies • Bond enthalpy is measured by determining how much energy is required to break the bond. • The energy required to break the bonds between two covalent bonded atoms is called bond dissociation energy. • Average bond enthalpies are positive, because bond breaking is an endothermic process.

  40. Bond Enthalpies & the Enthalpy of Reactions • To estimate ∆H for a reaction, we have to compare the bond enthalpies of bonds broken to the bond enthalpies of the new bonds formed. • ∆ Hrxn= D(Bond enthalpies of bonds broken) − D(bond enthalpies of bonds formed) • CH4(g) + Cl2(g) to CH3Cl(g) + HCl(g) • ∆ H = [D(C—H) + D(Cl—Cl)] − [D(C—Cl) + D(H—Cl)] = [(413 kJ) + (242 kJ)] − [(328 kJ) + (431 kJ)] = (655 kJ) − (759 kJ) = −104 kJ

  41. Bond Enthalpy Table

  42. References • Lecture: Theodore E. B., Eugene, H. L. H., Bruce E. B., Catherine M., Patrick W., (2011). Chemistry: The Central Science (12 Ed). Prentice Hall. USA. • Laboratory: Theodore E. B., John H. N., Kenneth C. K., Matthew S. (2011). Laboratory Experiments for Chemistry: The Central Science (12 Ed). Prentice Hall. USA. • Theodore E. B., (2011). Solutions to Exercises for Chemistry: The Central Science. Prentice Hall. USA. • John M., Robert C. F. (2010). Chemistry (4 Ed): Prentice Hall Companion Website. • Chemistry Online at • Chemistry and You at • Teachers Notes •

  43. END