Presentation is loading. Please wait.

Presentation is loading. Please wait.

Valence electrons Electrons in the highest occupied energy level of an atom. These are the electrons that determine the element’s properties.

Similar presentations

Presentation on theme: "Valence electrons Electrons in the highest occupied energy level of an atom. These are the electrons that determine the element’s properties."— Presentation transcript:

1 Valence electrons Electrons in the highest occupied energy level of an atom. These are the electrons that determine the element’s properties.

2 Electron dot structures
Diagrams that show the valence electrons as dots. The core electrons and the nucleus are included in the symbol of the element

3 Octet Rule In forming compounds, atoms tend to achieve noble gas configuration. 8 electrons in the highest energy level.

4 IONS Atoms or groups of atoms that have a positive or negative charge.
Cations - positive ion resulting from loss of electrons. Anions - negative ions resulting from gain of electrons.

5 Formula Unit Represents an ionic compound.
Lowest whole number ratio of ions in the compound.

6 Formula units Ionic compounds form as repeating links in a crystal matrix. Each cation is bound to each neighboring anion The ions are “locked in place” NaCl is the smallest ratio that indicates this matrix

7 Ionic Bond Bonds resulting from the electrostatic attraction between oppositely charged ions. In an ionic compound the net ionic charge is 0.

8 Ionic Compounds Metal + Non-metal Polyatomic cation+Non-metal
Metal+ Polyatomic anion Solid at room temperature High melting point >300°C

9 Ionic compounds NaCl Na2SO4
CaCO3 Crystalline solids that have high melting points. They are often soluble in water They conduct electricity when in solution, or when molten.

10 Predicting the formula
Identify the charge of the cation Identify the charge of the anion Make a “T” table Add ions until the positive charge equals the negative charge

11 Predicting the formula
Polyatomic ions are groups of atoms that stay together, they are treated like “super atoms” The entire group carries the charge. Table 9.3 p.257 NH4+

12 Lets practice !!!! Potassium + Phosphorus Lithium + Selenium
Aluminum + Chlorine Gallium + Sulfur Magnesium + Iodine Sodium + Carbonate Sodium + Hydrogen Carbonate Strontium + Phosphate Ammonium + Chromate Barium + Acetate

13 Lets practice !!!! Potassium + Phosphorus K3P Lithium + Selenium Li2Se
Aluminum + Chlorine AlCl3 Gallium + Sulfur Ga2S3 Magnesium + Iodine MgI2 Sodium + Carbonate Na2CO3 Sodium + Hydrogen Carbonate NaHCO3 Strontium + Phosphate Sr3(PO4)2 Ammonium + Chromate (NH4)2CrO4 Barium + Acetate Ba(C2H3O2)2

14 Lets Review!!! calcium chloride cesium oxide aluminum perchlorate
barium sulfide sodium dichromate aluminum phosphate calcium carbonate sodium carbonate

15 Lets Review!!! calcium chloride CaCl2 cesium oxide Cs2O
aluminum perchlorate Al (ClO4)3 barium sulfide BaS sodium dichromate Na2Cr2O7 aluminum phosphate Al PO4 calcium carbonate CaCO3 sodium carbonate Na2CO3

16 Compounds with transition metals
Transition metals can have more than one charge. You may have more than one possible compound: FeO, or Fe2O3 Make tables & work backwards to determine cation charge Indicate charge with a roman numeral REMEMBER THE TABLE MUST BE BALANCED !!!!! Fe+? O-2 iron(II) oxide Fe+? O-2 iron(III) oxide

17 Some Ions we need to just Know
Silver is always +1 Zinc is always +2 Cadmium is always +2 Do not use a roman numeral with these Iron may be +2 or +3 Tin may be +2 or +4 Lead may be +2 or +4 More in table 9.2 p.255

18 Lets Review!!! potassium oxide strontium nitride strontium nitrate
strontium nitrite aluminum hydroxide magnesium sulfate iron(III) oxide silver oxide

19 Lets Review!!! Potassium Oxide K2O Strontium Nitride Sr3N2
Strontium Nitrate Sr(NO3)2 Strontium Nitrite Sr(NO2)2 Aluminum Hydroxide Al(OH)3 Magnesium Sulfate MgSO4 Iron(III) Oxide Fe2O3 Silver Oxide Ag2O

20 Lets practice!! Na2S Hg2S Na2Cr2O7 Hg2Cr2O7 CuO

21 Lets Practice Answers Na2S sodium sulfide Hg2S mercury(I) sulfide
Na2Cr2O7 sodium dichromate Hg2Cr2O7 mercury(I) dichromate CuO copper(II) oxide

22 Lets Practice calcium carbonate ammonium sulfate copper(I) phosphate
chromium(IV) acetate cadmium perchlorate

23 Lets Practice calcium carbonate CaCO3 ammonium sulfate (NH4)2SO4
copper(I) phosphate Cu3PO4 chromium(IV) acetate Cr(C2H3O2)4 cadmium perchlorate Cd(ClO4)2

24 Links to practice tests and games
Interactive link Interactive link 2 Interactive link 3 Interactive link 4

25 Metallic Bonds The force of attraction that holds metals together. The attraction of the free floating electrons for the positively charged metal ions

26 Metallic Properties Malleable Ductile Conduct heat and electricity.

27 Single covalent bond A bond in which two atoms share a pair of electrons between them in order to achieve noble gas configuration.

28 Structural formulas Chemical formulas that show the arrangement of atoms in molecules and polyatomic ions. Each dash represents a pair of shared electrons.

29 Unshared pairs Pairs of valence electrons that are not involved in bonding, not shared between atoms. Also called lone pairs or non-bonding pairs

30 Double covalent bond Two atoms share two pairs of electrons between them to attain noble gas configuration. O2 and CO2

31 Triple covalent bond Two atoms share three pairs of electrons between them to attain noble gas configuration. N2

32 Coordinate covalent bond
A covalent bond in which one atom contributes both bonding electrons. CO and NH4+ and N2O

33 Exceptions to the octet rule
NO2 BF3 PCl5 SF6

34 Law of Definite Proportions
In samples of any chemical compound, the masses of the elements are always in the same proportions.

35 Law of Multiple Proportions
When two elements form more than one compound, the different masses of one element that combine with the same mass of the other element are in the ratio of small whole numbers.

36 Isoelectronic ions Ions containing the same numbers of electrons.
Generally for isoelectronic ions size decreases as nuclear charge increases.

37 Bond energy The energy required to break a bond. Table 8.4 p.365
Bond length- the distance where energy is minimum.Table 8.5 p.365

38 Coulomb’s Law E=2.31x10-19Jxnm Q1Q2 r E= energy in joules
r= distance between ion centers in nm Q1&Q2= ion charges

39 Molecular Compounds All non-metals covalently bonded.
Solid, liquid or gas Low melting point <300°C Smallest representative particle is a molecule.

40 Non-polar covalent bond
A covalent bond in which the electrons are shared equally. The two atoms have nearly the same electronegativities

41 Polar covalent bond A covalent bond in which the electrons are not shared equally. The more electronegative atom will pull more of the electrons toward itself.

42 Polar molecule One end of the molecule has a slightly positive charge and one has a slightly negative charge. This is called a dipole. Depends on the shape.

43 Lattice energy Lattice Energy=k(Q1Q2) r
The change in energy that takes place when separated gaseous ions are packed together to form an ionic solid. Lattice Energy=k(Q1Q2) r

44 Use the following to calculate DH°f of BaCl2(s).
Lattice energy= kJ/mol 1st ionization Ba= 503kJ/mol 2nd ionization Ba= 965kJ/mol Electron affinity Cl=-348kJ/mol Bond energy Cl2=239kJ/mol DH sublimation Ba=178kJ/mol

45 Bond Energies & Enthalpy
DH=SD(bonds broken)- SD(bonds formed) S=sum of terms D=bond energy per mol of bonds, always positive.

46 Localized Electron Bonding Model
A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. Lewis Structures VSEPR Theory

47 Writing Lewis Structures
Sum the valence electrons from all the atoms. Use a line to show a pair of electrons between each pair of bound atoms (Bonding Pairs)

48 Writing Lewis Structures
Arrange the remaining electrons to satisfy the duet rule for hydrogen and the octet rule for the 2nd row elements. (Lone Pairs) Double or triple bonds may be needed.

49 Comments on Octet Rule C,N,O,F obey octet rule. B and Be often have less than 8 electrons. Very reactive. 2nd row never exceed rule. 3rd row and up often obey octet rule but may exceed it., due to d orbitals.

50 When writing lewis structures satisfy the octet rule for the atoms first. Place any remaining electrons on the elements that have available d orbitals.

51 Resonance structures Occur when it is possible to have two or more valid electron dot structures for the same molecule or ion. SO3, SO2

52 Formal Charge (FC) A method to decide which of many possible non-equivalent Lewis structures is most likely to occur. Atoms in molecules try to achieve FC as close to 0 as possible.

53 FC=(# valence e- on free atom) -(# valence e- assigned to the atom in the molecule).
(Valence e-)assigned = (# lone pair electrons) + 1/2(#shared electrons)

54 VSEPR theory Valence-shell-electron-pair repulsion theory.
Because electron pairs repel molecular shape adjusts so the valence electron pairs are as far apart as possible.

55 Hybrid orbitals In hybridization several atomic orbitals mix to form the same number of equivalent hybrid orbitals

56 Sigma bonds Formed along the axis that joins the atomic nuclei when two atomic nuclei combine to form a molecular orbital.

57 Pi bond Electron in pi bonds are found in sausage shaped regions above and below, or in front and behind the bond axis.

58 Paramagnetic molecules
Show an attraction to an external magnetic field. Molecules contain one or more unpaired electrons.

59 Diamagnetic Molecules
Molecule is repelled by an external magnetic force. Associated with paired electrons

60 Electronegativity The ability of an atom in a compound to draw electrons to itself. Pauling electonegativity values Table 14.2 p.405 Large electronegativity differences correspond to ionic bonds

Download ppt "Valence electrons Electrons in the highest occupied energy level of an atom. These are the electrons that determine the element’s properties."

Similar presentations

Ads by Google