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Chemical Bonding CHAPTER SIX (6).  Explain, using your periodic table, how to calculate (find) the following in an atom:  Atomic number _____________________.

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Presentation on theme: "Chemical Bonding CHAPTER SIX (6).  Explain, using your periodic table, how to calculate (find) the following in an atom:  Atomic number _____________________."— Presentation transcript:

1 Chemical Bonding CHAPTER SIX (6)

2  Explain, using your periodic table, how to calculate (find) the following in an atom:  Atomic number _____________________  Atomic mass _______________________  Identity of the atom _________________  Number of protons __________________  Number of neutrons __________________  Number of electrons ________________ Daily Assignment

3  Remember that the desire of every element is to be “happy,” that is, to have a full valence shell.  Atoms are willing to do whatever it takes to accomplish this, including align themselves with some “seedy” atoms  Having a full valence shell is also known as having a stable electron configuration The Purpose of an Element

4  In order to talk about bonding, we need to be able to visualize the elements we’re talking about. Here’s how we do that: Drawing Elements

5  When atoms bond together to form compounds, they may gain or lose electrons. An atom that has gained or lost electrons is called an ION. atom Na gives an electron to Cl, creating two ions. Ionic Bonding

6  The atom Na has 11 (-) electrons, balanced by 11 (+) protons in the nucleus.  The atom Cl has 17 (-) electrons, balanced by 17 (+) protons in its nucleus. Let’s Look Closer…

7  The octet rule says that atoms tend to gain, lose or share electrons so as to have eight electrons in their outer electron shell (HAPPY).  Na atom Cl atom Na has1 e - Cl has 7e - ( it can lose one or gain seven) (it can lose seven or gain one) If Na gives up its electron to Cl, both will have 8e -. THE OCTET RULE

8  When a bond is formed, Na gives up one electron to Cl, reducing its number of electrons to 16 and increasing Cl’s electrons to 18.  Na forms a 1+ charge, since it now has one more proton than electron. Cl forms a 1- charge, since it now has one more electron than proton.  The charge is written as a superscript number to the right of the element’s symbol.

9 Mg loses two electrons to form ______________ Oxygen gains two electrons to form __________ Phosphorus gains three electrons to form _______ Carbon gains four electrons to form _________ Calcium loses two electrons to form __________ How was N 3- formed? ___________________ How was Ag 1+ formed? ___________________

10 Why do most atoms tend to form compounds, as opposed to existing in their unbonded state?

11 To find the number of protons, neutrons, and electrons in an ion, you need to look at the atomic number, mass, and the charge on the ion. Remember, atomic number = # of protons atomic mass = protons + neutrons neutrons = atomic mass – atomic number Now, to calculate the number of negative electrons, you must balance the charge of the ions with the number of positive protons.

12  Atomic number = ______  Atomic mass = _______  Charge = ___________  Element name _______  Element symbol _______  # protons __26___  # neutrons __30__  # electrons ___23__  Atomic number = __26__  Atomic mass = _______  Charge = ___________  Element name _______  Element symbol _______  # protons _____  # neutrons __29__  # electrons __24___

13  Atomic number = __26__  Atomic mass = __56___  Charge = ___+3_____  Element name __Iron  Element symbol __Fe___  # protons __26___  # neutrons __30__  # electrons ___23__  Atomic number = __26__  Atomic mass = __55___  Charge = ___+2____  Element name __Iron__  Element symbol __Fe___  # protons __26_  # neutrons __29__  # electrons __24___

14  Atomic number = __16__  Atomic mass = _______  Charge = _____2-____  Element name ______  Element symbol ____  # protons ______  # neutrons __18__  # electrons _____  Atomic number = _____  Atomic mass = ___7____  Charge = ___________  Element name _______  Element symbol _______  # protons __3___  # neutrons _____  # electrons ___2__

15  Atomic number = __16__  Atomic mass = __34__  Charge = _____2-____  Element name ___Sulfur_  Element symbol __S___  # protons __16___  # neutrons __18__  # electrons __18___  Atomic number = __3___  Atomic mass = ___7____  Charge = ____+1____  Element name __Lithium_____  Element symbol __Li___  # protons __3___  # neutrons __4__  # electrons ___2__

16  Atomic number = ______  Atomic mass = __40__  Charge = _____1+____  Element name _______  Element symbol ___K___  # protons _____  # neutrons ____  # electrons _____  Atomic number = __19__  Atomic mass = _______  Charge = ___________  Element name _______  Element symbol _______  # protons _____  # neutrons __22__  # electrons __19___

17  Atomic number = __19__  Atomic mass = __40__  Charge = _____1+____  Element name __Potassium___  Element symbol ___K___  # protons __19__  # neutrons __21_  # electrons __18__  Atomic number = __19__  Atomic mass = ___41____  Charge = _____0______  Element name ___Potassium____  Element symbol ___K___  # protons ___19__  # neutrons __22__  # electrons __19___ Periodic Table

18 Explain the “Octet Rule” _______________________________________ _______________________________________ _______________________________________ ________  One atom has 2 electrons in its outer shell. Another atom has six. Will they combine to form a compound?  What about an atom with two electrons and 2 atoms with seven outer shell electrons?

19  Atoms will combine with other atoms to achieve a chemically stable arrangement of their electrons.  Recall, the OCTET RULE drives chemical bonding. A chemically stable arrangement of electrons is attainable by completely filling the outer shell (with eight electrons*).

20 Atoms, like small children, will always do what benefits them most. In this case, their electrons can either be shared between atoms or transferred from one atom to another to form a chemical bond.

21 Ionic Bonds occur when one atom loses electrons to another atom that gains them. Ions arrange themselves into crystal lattices.

22 Ionic Bonds = transfer of electrons between atoms in a reaction * Only outer shell electrons shown *

23  To write the newly formed ionic compound, take the charges after the reaction, criss-cross them, and drop the sign.  NaClMgS KI

24  Draw a dot diagram and complete the charges for the reaction of:  Sodium Boronate

25 Electrons are “shared” between atoms. Since they spend time in the electron clouds of each atom, the atom “feels” complete and happy.

26  The fluorine atoms, being nice to each other as they are, decide to share two electrons - one from each of them. This obviously means that one of the atoms - the one donating - has a complete outer shell; the other atom has a gap for two electrons, e.g. something with six electrons in its outer shell).  Despite the fact that each atom has its own electron cloud, each atom also feels like it has a full outer shell because of the shared electrons.  All electrons are still spinning around each atom, but because of the sharing, the atoms are bonded together in order to maintain the sense of a full outer shell.

27  Most atoms become chemically stable by sharing electrons. This does not result in the formation of ions (the atoms do not have charges). These electrons belong to both H and O. At any point in time they can exist in either electron cloud.

28  Do the following reactions at your seat. Draw the dot diagram. Are they covalent or ionic bonds?  Hydrogen and Hydrogen  Chlorine and Chlorine  Hydrogen and Chlorine

29 Copy the reactions listed above. Determine the type of chemical bond formed during the reactions.

30 Sometimes elements do not combine with other elements. Instead, atoms of the same element combine. When this happens, we write it a special way. Take Nitrogen, for example. Single bond Double BondTriple Bond

31 Those elements on the right of the periodic table crave electrons more than others, and those at the top crave them more than those at the bottom.

32 Occurs when atoms get greedy and “hog” the shared elements. Example: Water

33  Not all molecules that contain polar bonds are polar molecules. It depends on how the electrons are shared.  Water is a polar molecule because the oxygen has more need for the electrons.  Carbon Dioxide shares the electrons between the carbon and oxygen atoms

34  The attractive forces between polar molecules are greater than those in non polar molecules.  The positive and negative ends of a polar molecule are attracted to eachother and create bonding forces.  Water is an example

35  A chemical formula is a combination of element symbols and subscript numbers that is used to show the composition of a compound. MgCl 2 The formula above tells you that the compound magnesium chloride contains one atom of magnesium bonded to two atoms of chlorine. There are three total atoms in the compound.

36 List the name of each element, the number of atoms of that element, and the total number of atoms in the chemical formula. H 2 O MgF 2 NH 4 I

37  Parentheses are used around polyatomic ions in a chemical formula. The subscript number to the right of the parentheses is used as a multiplier for the atoms inside. Ba(OH) 2 (NH 4 ) 2 S Au 2 (CO 3 ) 2

38 Coefficients are used in a chemical formula to represent the number of units of the compound present. They act as a multiplier for all atoms present in the formula. 3 H 2 O 4 Ba 3 (PO 4 ) 2

39 List the number of atoms of each element and the total number of atoms in the compound for the formulas listed below. 2 C 4 H 6 (NH 4 ) 2 CO 3 5 Al 2 (SO 4 ) 3

40  You have learned that elements chemically bond together (ionic or covalent) to form new substances.  You have learned that a chemical formula indicates the number of atoms of each element contained in a chemical compound. magnesium chloride MgCl 2  Write a formula for a compound with two aluminum and three sulfur atoms. Write a formula for a compound with one beryllium and two iodine atoms.

41  Ionic compounds are formed from the transfer of electrons, resulting in the creation of a positive ion and a negative ion.  The compound is always written positive ion (cation) then negative ion (anion).  The negative ion (if it is an element) will have an –ide ending. calcium fluoride

42  Some metals form multiple ions. Alkalis, Alkaline Earths, and Aluminum form ions with positive charges equal to their group number, such as K +, Ca 2+, and Al 3+  Many transition metals form more than one type of ion. Examples include copper I and copper II. Roman numerals designate the charge.  Cu 2 OCuO

43  Compounds that join together in a regular way and act as one unit. It can have either a negative or positive charge.  Some examples of polyatomic ions: AmmoniumNH 4 + AcetateC 2 H 3 O 2 - HydroxideOH - PeroxideO 2 2- NitrateNO 3 - Hydrogen Sulfate SulfateSO 4 2- HSO 4 - CarbonateCO 3 2- Hydrogen Carbonate PhosphatePO 4 3- HCO 3 - ChromateCrO 4 2- Hydrogen Phosphate SilicateSiO 3 2- HPO 4 2- PermanganateMnO 4 - Hypochlorate OCl -

44 1. Write the symbol for the positive and negative ions. 2. List the oxidation number (charge) for each. This can be found on your periodic table or listed in the formula name, following the positive ion as a roman numeral. Find the oxidation numbers for the following ions: AgAu (I)Pb (IV)S BePb (II)AlCa H ClOF

45  Balance the charges by the criss-cross method. Criss- cross and drop the charge!  Place all parentheses around all polyatomic ions. Make sure the criss-crossed number is placed outside the parentheses.  Reduce subscript numbers to reflect the Least Common Multiple.

46 Sodium nitride Magnesium bromide Barium sulfide Lead (IV) oxide Potassium selenide Chromium nitride Copper (I) oxide Copper (II) oxide

47 Write the formula for barium chloride. Explain, using electron dot diagrams, how this compound forms and whether it will be an ionic or covalent bond formed. (I’ll start you off) Ba + Cl + Cl

48 PPlace all parentheses around all polyatomic ions. Make sure the criss-crossed number is placed outside the parentheses. RReduce only the subscript numbers to the outside of the parentheses. Lead (II) phosphatemagnesium acetate Ammonium sulfidesilver silicate

49  Aluminum SulfateMagnesium Acetate  Hydrogen PeroxideAmmonium Nitride  Beryllium PhosphateCalcium Hydrogen Phosphate  Potassium ChromatePotassium Hydroxide

50 Mercury (II) carbonatePlatinum (I) sulfide Platinum (IV) sulfateammonium nitrate Calcium iodidehydrogen peroxide

51 As you recall, elements undergoing molecular (covalent) bonding are sharing electrons, so no charges (or ions) are formed during the formation of the compound. Molecular compounds are named and written using a system of Greek Prefixes, which must be memorized.

52 One = mono Six = hexa Two = di Seven = hepta Three = triEight = octa Four = tetraNine = nona Five = pentaTen = deca

53  Just let the prefix for the element in the formula become the subscript number for that element. If there is no prefix, the subscript is 1. Examples: dihydrogen monoxide trisilicon pentabromide tetraphosphorus hexasulfide carbon dioxide arsenic trichloridedecaphosphorus heptoxide

54  1. disulfur hexoxide  2. ammonium sulfide  3. iron (III) carbonate  4. iron (II) carbonate  5. aluminum phosphite  6. gold (III) nitride  7. diphosphorus pentabromide  8. silicon sulfide  9. tetraradon heptafluoride  10. potassium oxide

55  Write the name of the first element  Write the name of the second element changing the name to an _ide ending.  Reverse cris-cross the subscript numbers to determine the charge of both ions. Add charges to periodic table

56  Check the periodic table to determine the number of positive charges that the first element forms.  If it only forms one, and it is the one listed, then you are done  If it only forms one, but it is not the one listed, then it has been reduced. Determine the multiplier and multiply both charges to determine the correct charge formed.  If it forms more than one, write the Roman numeral equivalent of that charge in the compound’s name.  If it forms more than one, yet none are listed, it has been reduced. See above, and use the Roman numeral.

57 AlBr 3 FeCl 3 NiBr 2 GeO 2

58 Ni 2 O 3 ZnCl 2 PtS 2 Cu 3 P 2

59 Follow all of the same rules, with the following exceptions:  Never change the name of a polyatomic ion  Place parentheses around polyatomic ions before reverse criss- crossing (if they are not already there)  Check the charges of element and polyatomic ions to check for a reduction.

60 Au 2 (CO 3 ) 3 FeSO 4 Ca 3 (PO 4 ) 2 (NH 4 ) 2 CrO 4

61 Ba(C 2 H 3 O 2 ) 2 Pt(NO 3 ) 4 Ni 2 (SiO 3 ) 3 HgSO 3

62 Simply let the subscript number become the prefix for the element’s name.  The first element keeps its name. It only gets a prefix if it has a subscript in the formula  The second element always gets an – ide suffix (ending). The second element always gets a prefix. N 2 0P 2 O 5 S 3 Cl 6 CO

63  Metals have special bonds because of the way they configure  Metals form regular patterns and share electrons with all the atoms around it, effectively forming cations  This forms a crystalline lattice like those in polar bonds  The metal is effectively neutral, though.

64  Properties of Metals explained ◦ Conductivity  Metals are good conductors of electricity, but why?  The lattice structure of metal atoms allows electrons to flow freely through the metal ◦ Malleability  A block of ice, if hit with a hammer, would shatter because of its ionic bonds  Although strong, the lattice in a metal is not as rigid as the lattice in ionic bonds. This allows the atoms to move but not break because the electrons keep the metal anions in place

65  An alloy is a mixture of metals to give it more desirable properties, such as hardness or resistance to rust  Copper Alloys ◦ First alloy was Bronze – a mixture of copper and tin ◦ Brass – another alloy of copper, formed by mixing copper and zinc

66  Iron Alloys ◦ The best known iron alloy is steel ◦ Steel is a combination of iron and carbon, which hardens the iron. ◦ Stainless steel contains about 10% of chromium, which makes it shiny and resistant to rusting ◦ Other elements, such as phosphorus, manganese, sulfur, and silicon, are added to steel to make strong, flexible metals used in bridges

67  Aluminum alloys are used in making airplane sheet metal (mixed with copper or manganese)  Aluminum is also mixed with magnesium to create a lightweight, but strong metal (sometimes used in performance bicycle frames)


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