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Unit 3 Part 3

Unit 3 Part 3. Pg 9-15. Determining the Charge (Transition metals). Determining charge when given the chemical formula: Work backwards: What is the charge on iron? FeCl 3

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Unit 3 Part 3

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  1. Unit 3 Part 3 Pg 9-15

  2. Determining the Charge (Transition metals) • Determining charge when given the chemical formula: • Work backwards: • What is the charge on iron? FeCl3 • We know chlorine has a charge of -1 and there are 3 chlorine ions, so copper must have a charge of +3 in order for the compound to have a neutral charge.

  3. Determining the Charge (Transition metals) • What is the charge on copper? Cu(NO3)2 • We know NO3 has a charge of ______ (from our common ion sheet), so copper must have a charge of ______. This is written as copper (II).

  4. Determining the Charge (Transition metals) • What is the charge on Iron in Fe2O3? • Oxygen has a charge of ______ normally. There are ______ oxygens so the oxygens have a total charge of ______ in total. This means that the irons must have a total charge or ______. There are 2 iron atoms and the molecule is neutral so they must each have a charge of ______ in order to equal ______.

  5. Determining the Charge (Transition metals) • What is the charge on iron in FeCO3? • First we need to find the polyatomic ion (it won’t always be in brackets). A compound contains a polyatomic ion if it is ionic and there are more than ______ elements. The polyatomic ion will be the first few elements if it is a _______________________ and the last few elements if it is an _______________________. Most polyatomic ions are anions, so usually the polyatomic ion will be at the _______________________. • Carbonate has a charge of ______ so iron has a charge of ______.

  6. Determining Charge Examples: • PbCl2

  7. Determining Charge Examples: • As(OH)5

  8. Determining Charge Examples: • FePO4

  9. Covalent Molecules • Number of Bonds: How to determine the number of covalent bonds that form: • Determining the number of covalent bonds that form is similar to determining charge. We look at how many electrons an atom needs to fills its octet (this is the number of bonds that will form). When atoms bond covalently they share their valence electrons.

  10. Number of Bonds • Carbon has _____ valence electrons. It needs _____ electrons to fill its valence shell. This means it will form _____ bonds (if they are single, or it could form 2 double bonds, or 1 triple bond and one single bond).

  11. Number of Bonds • Oxygen has 6 valence electrons and needs _____ to fill its valence shell. This means it will form _____ bonds (if they are single, or 1 double bond).

  12. Number of Bonds • Nitrogen has _____valence electrons and needs _____ electrons to fill its valence shell. This means it will form _____bonds (if they are single).

  13. Lewis Dot (Covalent) • Draw Lewis dot structure for molecules (covalent): • Fluorine bonding with itself (F2): • Determine how many atoms of each element there will be (if it’s not given). • Draw the lewis dot for each atom. (In the space below) • Look at the lone electrons. These are the electrons involved in bonding. Circle one lone electron from one atom and one lone electron from another. This will form a bond. If there are three or more atoms, make sure every atom is involved in the bonding. • Redraw the molecule so that the atoms are beside one another and the shared electrons are between them. • If asked to draw a structural diagram, draw lines to replace bonded pairs of electrons.

  14. Examples: • Oxygen bonding with itself: • Chlorine bonding with hydrogen:

  15. Examples • CF4: • Nitrogen bonding with itself:

  16. VSEPR HANGMAN What does VSEPR stand for? __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ ABCDEFGHIJKLMNOPQRSTUVWXYZ

  17. VSEPR • What do the molecules really look like? • Next we will look at what a molecule looks like in 3 dimensional space. We use molecular geometry and VSEPR Theory. VSEPR stands for Valence Shell Electron Pair Repulsion. The VSEPR theory determines the shape of a molecule by looking at the electrons surrounding the central atom and whether they are shared pairs of lone pairs (bonding or non-bonding pairs).

  18. VSEPR • Determine Molecular geometry using VSEPR: • Determine the Lewis dot formula • Determine the total number of electron pairs around the central atom • Use the table provided to determine the electron pair geometry • Use the table provided to determine the shape • Use the diagram chart to draw a 3D diagram

  19. VSEPR Geometry 3D diagrams: E = Central Atom X =Bonded Atom = Unbonded pair of electrons

  20. Examples: • OF2 • Oxygen has 4 pairs of electrons around it, _____ are bonded _____ are lone, so it has a geometry of ______________________ ______________________.

  21. The 3D drawing for tetrahedral bent is: • So you just need to fill in the atoms where they belong: 

  22. PCl3 PCl3 has _____ pairs of electrons around it, _____ of the pairs are bonded. This means that PCl3 has a geometry of _____________________, ___________________ _____________________.

  23. Sometimes in a double or triple bond you have to assume that the atom has only one shared pair of electrons in order to get the proper geometry. CS2 Even though this actually has 4 shared pairs of electrons and all are bonding electrons, the shape is linear. In this molecule we assume that Carbon has 2 pairs of electrons and both are bonded.

  24. Exceptions to the octet rule • Normally assume that atoms want to have 8 valence electrons, however sometimes atoms like to have more. In situations like this we say that the atom breaks the octet rule. In this class you will always be told if an element breaks the octet rule.

  25. XeF4 • Xe has 8 valence electron and Fluorine has 7, http://www.polleverywhere.com/multiple_choice_polls/MTcxMTg5MzUxOQ • This molecule has ______ pairs of electrons around it, and 4 of them are involved in bonds so its geometry is ______________________, ______________________ ______________________. • 3D drawing:

  26. Intermolecular Forces • Inter is latin for “between” • Occurs when bonds or forces occur between molecules (covalent). • Intermolecular forces hold a molecule in a certain spot. • The three types of intermolecular forces we will look at are polar bonds, hydrogen bonds and van der waals forces.

  27. Polarity • Non-polar molecules have equal, or approximately equal pull on the electrons. This occurs when the atoms in the molecule are the same or have a similar electronegativity. Electronegativity generally increases as you move from left to right in a period and decreases as you move from the top to the bottom of a family.

  28. Polarity • Polar molecules have a slightly positive or slightly negative charge due to one atom having a stronger pull on the electrons. This occurs when there is a large difference in electronegativity so one atom pulls stronger on the electrons than the other atom. The more electronegative atom will have a partial negative charge because it has a stronger pull on the electrons, and the less electronegative atom will have a partial positive charge because the electrons are being pulled away.

  29. Polarity • Example, Chlorine has an electronegativity of 3.16 and hydrogen has an electronegativity of 2.0. Since chlorine has a stronger pull on the electrons they will be closer to the chlorine atom and the chlorine atom will have a slightly negative charge. This means the hydrogen will have a slightly positive charge.

  30. Polarity • To show the direction of polarity we draw a dipole. A dipole is drawn to show where the electrons in the molecule are more likely to be found. It is drawn like this: H Cl

  31. Polarity • Another way to show the slight negative charge is to use the symbol that stands for slightly.

  32. Polarity • The difference in the electronegativity can tell us what type of bond is formed.

  33. Polarity • If the difference is 0-0.4 then the bond is non-polarcovalent

  34. Polarity • If the difference is 0.4- 1.7 then the bong is polar covalent

  35. Polarity • If the difference in electronegativity is greater than 1.7, then the bond is ionic. If the difference in electronegativity is very very large then it is usually an ionic bond because one atom pulls so hard on the electrons that it pulls them completely away from the other atom.

  36. Polarity • Remember that there is no distinct line of separation so the rules above are more like guidelines. Remember an ionic bond will occur between a metal and a non-metal and covalent is between 2 non metals.

  37. Example • H-I

  38. Assignment • (See Polar vs. Non-Polar Assign Pg 42)

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