* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Download Lone pairs
History of chemistry wikipedia , lookup
Computational chemistry wikipedia , lookup
Physical organic chemistry wikipedia , lookup
Electrical resistivity and conductivity wikipedia , lookup
Low-energy electron diffraction wikipedia , lookup
Halogen bond wikipedia , lookup
Photoelectric effect wikipedia , lookup
Atomic nucleus wikipedia , lookup
Hydrogen bond wikipedia , lookup
Molecular orbital wikipedia , lookup
Gaseous detection device wikipedia , lookup
X-ray photoelectron spectroscopy wikipedia , lookup
Auger electron spectroscopy wikipedia , lookup
Metastable inner-shell molecular state wikipedia , lookup
Aromaticity wikipedia , lookup
Rutherford backscattering spectrometry wikipedia , lookup
Light-dependent reactions wikipedia , lookup
Atomic orbital wikipedia , lookup
Photosynthetic reaction centre wikipedia , lookup
Bent's rule wikipedia , lookup
Electronegativity wikipedia , lookup
Resonance (chemistry) wikipedia , lookup
Atomic theory wikipedia , lookup
Molecular orbital diagram wikipedia , lookup
Bond valence method wikipedia , lookup
History of molecular theory wikipedia , lookup
Electron configuration wikipedia , lookup
Metallic bonding wikipedia , lookup
Chemical Bonding: Bonding Theory and Lewis Formulas Periodicity There are certain trends within the periodic table which affect reactivity and the ability to form bonds. Periodic Law When elements are arranged in order of increasing atomic #, elements with similar properties appear at regular intervals. 250 Atomic Radius (pm) Cl 200 F 150 Br 100 50 0 0 5 10 Atomic Numbe r 15 20 Chemical Reactivity Families 1 2 3 4 5 6 7 Similar valence e- within a group result in similar chemical properties Chemical Reactivity 1 2 3 4 5 6 7 Alkali Metals Alkaline Earth Metals Transition Metals Halogens Noble Gases Properties that affect Reactivity Atomic Radius size of atom Atomic Radius 250 K Atomic Radius (pm) Halogens 200 Na Li 150 100 Ar Ne 50 Noble Gases 0 0 5 10 Atomic Number 15 20 Atomic Radius Atomic radius increase as you move top to bottom. Atomic radius increase as you move right to left. 1 2 3 4 5 6 7 Ionization Energy Again Ionization whenenergy you compare is the amount the ionization of energy energy neededoftoelements turn a stable thereelement is a trend.. into an ion. He 1st Ionization Energy (kJ) 2500 Ne 2000 Ar 1500 1000 500 Li Na K 0 0 5 10 Atomic Number 15 20 Ionization Energy Ionization energy increases as you move right and down. 1 2 3 4 5 6 7 Ionization Energy Why are atomic radius and ionization energy opposite? In small atoms, the negative electrons (e-) are closer to the positive nucleus where the attraction is stronger. 50 mμ 20 mμ Small Atoms Large Atoms Valence Electrons, Valence Energy Levels & Valence Orbitals Periodic trends are related to the number of valence electrons an element has. Valance electrons are those electrons occupying the highest energy level of an atom (the outside shell). How Many Valence Electrons? First Level: 2eSecond Level: 8eThird Level: 8e- ***You can tell how many valence electrons are in each orbital by counting the number of elements on each row of the periodic table.*** Ex. The magnesium atom has 12 protons and 12 electrons. The maximum number of electrons in the first energy level is 2 e- That leave 10 electrons left to place. The next energy level cam only hold 8 electrons. That only leaves 2 electrons that are TRUE valence electrons Orbital: a region in space in which an electron with a given energy is likely to be found. There are four valence orbitals within the valence energy level of an atom (1s and 3p’s) There are few exceptions to this rule Hydrogen Helium Electrons will occupy all valence orbitals before forming electron pairs. Empty bus seat rule Normally a maximum of 8 electrons may occupy a valence energy level. This is known as the octet rule. urinaltest 2 electrons in first energy level. You place the other 4 electrons in each of the four orbital. \Leaving you with 4 more electrons to place. Start at the top and go clockwise. Electron Dot Diagrams (Lewis Symbols) Electron dot diagrams can represent atoms (neutral) or ions (charged). ONLY show the atom’s valence electrons! – These are the only electrons involved in a chemical reaction (the electrons in the outer most ring!) Dot Diagram Trends Follow these steps: 1) Write the atomic symbol for the atom. This symbol represents the nucleus and the core electrons that do not participate in the chemical bonding. 2) Dots () represent the electrons in the valence energy level of the atom. Arrange these dots around the atomic symbol. 3) One dot must be placed in each of the four orbitals before any electron pairing occurs. 4) Begin with the fifth electron to make lone pairs. (if you have to) 5) There is a maximum of 8 electrons that can be drawn. Lets try some… Calcium Ca Oxygen O Bromine Carbon Br C Bonding Electrons versus Lone Pairs Bonding electrons are unpaired electrons that are involved in bond formation. Paired electrons are called lone pairs and are generally not involved in bond formation. Bond Types There are 3 types of bonds that can be formed These are determined by which elements combine Types of Bonds 1. IONIC 2. COVALENT Bond Formation e- are transferred from metal to nonmetal e- are shared between two nonmetals Type of Structure crystal lattice true molecules Physical State solid liquid or gas Melting Point high low Solubility in Water yes usually not Electrical Conductivity yes (solution or liquid) no Other Properties odorous Metallic Bonding 2 metals share electrons but no chemical reaction occurs Valence electrons are free to move about between the atoms Positive ions surrounded by a “sea” of mobile electrons Allows metals to be formed into any shape 1. Ionic Bonding A complete transfer of electrons occurs in an ionic bond. Valence Electrons 1. Ionic Bonding 1 7 8 8 2. Covalent Bonds 2. Covalent Bonding Results in a mutual sharing of electrons between the two non-metals. This Sharing can be: Equal = nonpolar covalent Unequal = polar covalent Nonpolar Covalent Bond e- are shared equally between both nucleus. Electron “Cloud” is symmetrical. Polar Covalent Bond e- are shared unequally asymmetrical e- density results in partial charges (dipole) + Polar Covalent Bond Example: H2O(l) H + O H 3.4 - 2.2 = 1.2 POLAR Bond Polarity Most bonds are a blend of ionic and covalent characteristics. Difference in electronegativity determines bond type. Electronegativity Electronegativity: The attraction an atom has for a shared pair of electrons….aka….the strength an atom has to hold onto or take electrons. Trends in Electronegativity Name That Bond!!!! Electronegativity Difference Between Two Atoms. Type of Bond Between the Atoms 1.7 Ionic 0.4 < 1.7 Polar covalent < 0.4 Nonpolar covalent Descriptions of Electrons in the Bond Transfer of electrons between metal and nonmetal Electrons shared unequally between unlike atoms Electrons shared equally between identical Bond Polarity Non-Polar (+) Ionic (-) Polar Bond Polarity Examples: Cl2 3.2-3.2=0.0 Nonpolar HCl 2.2-3.2=1.0 Polar NaCl 0.9-3.2=2.3 Ionic Covalent bonds are classified as: 1. single (sigma bond), 2. double (1sigma and 1pi bond), or 3. triple bonds (1sigma and 2 pi bonds) depending on the number of electrons shared between the two nuclei. Electron Dot Diagrams for Ionic Compounds Electrons are transferred from the metal to the nonmetal. results in a net negative charge for the nonmetal. And a net positive charge for the metal. Electron Dot Diagrams for Ionic Compounds Drawing Na Cl Drawing Mg Cl Cl [Mg]2+ -[Na]+ [ [Cl ] Cl ] Lewis Dot Diagrams For Molecular Compounds Octet Rule Remember… Most atoms form bonds in order to have 8 valence electrons. Octet Rule Exceptions: Hydrogen 2 valence e- Helium 2 Valence e- HOH Drawing Lewis Diagrams 1. Find total # of valence e-. 2. Arrange atoms - singular atom is usually in the middle. 3. Form bonds between atoms (2 e-). 4. Distribute remaining e- to give each atom an octet (recall exceptions). 5. If there aren’t enough e- to go around, form double or triple bonds. Drawing Lewis Diagrams CF4 1 C × 4e- = 4e4 F × 7e- = 28e32e- 8e24e- How many bonds can Carbon (C) form? F F C F F Drawing Lewis Diagrams BeCl2 1 Be × 2e- = 2e2 Cl × 7e- = 14e16e- 4e12e How many bonds can Beryllium (Be) form? Cl Be Cl Drawing Lewis Diagrams CO2 1 C × 4e- = 4e2 O × 6e- = 12e16e- 4e12e How many bonds can Carbon (C) form? O C O Polyatomic Ions To find total # of valence e-: Add 1e- for each negative charge. Subtract 1e- for each positive charge. Place brackets around the ion and label the charge. Polyatomic Ions ClO41 Cl × 7e- = 7e4 O × 6e- = 24e31e+ 1e32e- 8e24e- How many bonds can Carbon (C) form? O O Cl O O Polyatomic Ions NH4+ 1 N × 5e- = 5e4 H × 1e- = 4e9e- 1e8e- 6e2e- How many bonds can Nitrogen (N) form? H H N H H Resonance Structures Molecules that can’t be correctly represented by a single Lewis diagram. Actual structure is an average of all the possibilities. Show possible structures separated by a doubleheaded arrow. 1. 2. Resonance Structures SO3 O O S O 1. 3. 2. O O S O O O S O Don’t Panic!!!! Lewis dot diagrams can take some time to figure out. If you cannot immediately determine the right configuration, take a deep breath and try again. Think of these as puzzles, keep on working with the pieces until they fit together You Want What? So what is the difference between a Lewis dot diagram? A chemical formula? And a structural diagram? CH4 Lewis Dot Diagram ___________________ Chemical Formula ___________________ Structural Diagram ___________________ 3.3 Molecular Shapes and Dipoles I II III VSEPR Theory Valence Shell Electron Pair Repulsion Theory Stereochemistry: The study of 3-D spatial configuration Electron pairs (bond pairs and lone pairs) orient themselves in order to minimize repulsive forces. VSEPR Theory In non science speak: The interaction of electrons is what forms bonds. Electrons have a negative charge, so they want to be AS FAR AS POSSIBLE AWAY FROM EACH OTHER. VSEPR Theory Types of e- Pairs Bonding pairs - form bonds Lone pairs - nonbonding e- Lone pairs repel more strongly than bonding pairs!!! VSEPR Theory Lone pairs reduce the bond angle between atoms. Bond Angle Determining Molecular Shape Draw the Lewis Diagram. Count how many things attached to central atom double/triple bonds = ONE Lone pairs = ONE Shape is determined by the # of bonding pairs and lone pairs. Know the base shapes & their bond angles as well as the derived shapes Base vs. Derived shapes Base Shapes: AKA electron pair geometry. Include all lone pairs as a bond Derived Shapes: AKA molecular geometry. Only worry about actual bonds. 1. Base Shape: Linear (AX2) 2 total electron pairs 2 bond pairs 0 lone pairs 180o bond angle BeH2 2. Base Shape: Trigonal Planar (AX3) 3 total electron pairs 3 bond pairs 0 lone pairs 120o bond angle BH3 3. Derived Shape: Bent/Angular (AX2E1) 3 total electron pairs 2 bond pairs 1 lone pair < 120o bond angle SO2 4. Base Shape: Tetrahedral (AX4) 4 total electron pairs 4 bond pairs 0 lone pairs 109.5o Bond angle CH4 5. Derived Shape: Trigonal Pyramidal (AX3E1) 4 total electron pairs 3 bond pairs 1 lone pair < 109.5o bond angle NH3 6. Derived Shape: Bent/angular/v-shaped (AX2E2) 4 total electron pairs 2 bond pairs 2 lone pairs <109.5o bond angle H2O 7. Derived Shape: Linear (AX1E3) 4 electron pairs 1 bond pair 3 lone pairs 180o bond angle HF 8. Derived Shape: Linear (AX2E3) 5 total electron pairs 2 bond pairs 3 lone pairs 180o bond angles XeF2 Examples PF3 4 total 3 bond 1 lone AX3E F P F F TRIGONAL PYRAMIDAL <109.5° Examples CO2 2 total 2 bond 0 lone AX2 O C O LINEAR 180° Examples (You Try) NO2- C 2H 2 CH3OH Molecular Polarity: Dipole Theory Polar Molecule: The negative (electron charge) is not distributed symmetrically among the atoms Dipole Moment Direction of the polar bond in a molecule. Arrow points toward the more e-neg atom. + H Cl Determining Molecular Polarity Depends on: dipole moments molecular shape Nonpolar Molecules Dipole moments are symmetrical and cancel each other out. F Non-Polar BF3 B F F Polar Molecules Dipole moments are not symmetrical. The bonds are between elements with an electronegativity difference of between 0.3-1.7. 3.4 ***Presence of lone pairs is a good trick to decide if its polar***** H2O O Polar 2.2 H 2.2 + H Group Activity VSEPR shape and Polarity Flashcards. In groups of three. Find yourself a place to work in the room with a flat surface. Decide who is A, B, and C. Group Activity VSEPR shape and Polarity Flashcards. A’s Come up to the front and get three markers of different colours. B’s Come up to the front and get 5 coloured flashcards. C’s You just sit tight…for now…..your time is coming. Group Activity VSEPR shape and Polarity Flashcards. When all the group members are back and the group has all the materials you may do the following: A’s: On the back (side without lines) of two cards you will draw the base shapes of Linear & Trigonal Planar. Group Activity VSEPR shape and Polarity Flashcards. B’s: On the back (side without lines) of three cards you will draw the base shapes of Bent, derived shape of trigonal planar and Tetrahedral. C’s: On the back (side without lines) of three separate cards you will draw the two derived shapes of Tetrahedral. A,B & C: You’re the “Editors” to make sure their all correct pictures. Group Activity VSEPR shape and Polarity Flashcards. LINEAR 180° BENT <109.5° BENT <120° TRIGONAL PLANAR 120° TETRAHEDRAL <109.5° TRIGONAL PYRAMIDAL <109.5° You have about 10 minutes. Group Activity VSEPR shape and Polarity Flashcards. Once all three group members have finished drawing the base and derived molecules shapes on the flashcards…. Each group member will find a actual compound that is a match to the shapes they drew earlier Each person will then draw the structural formula of the compound they found on the other side of the flashcard, along with the chemical formula of the compound. You have 10 minutes to complete pictures. Group Activity VSEPR shape and Polarity Flashcards EXAMPLE FRONT: TETRAHEDRAL <109.5° BACK: CH4 VSEPR LAB CH2Cl2 CCl4 CH3Cl BH3 C2H2 PCl3 NF3 AlCl3 HClO C2Cl2 C2H4 C2H6 C2H2Cl NH3 HCN SiF4 Cl2O H2O2 Group Activity VSEPR shape and Polarity Flashcards. Once every member of your group has completed their flashcards, share them in your group, everyone must be able to do ALL of the shapes by themelves!! HINT: There are assignments in your assignment bundle that are very close to what you are doing….. You have about 10 minutes to share in your group. Meniscus? Capillary action? Is water attracted to a charged object? Intramolecular bonds: attractions within a molecule. Intermolecular forces: The weak forces or bonds between molecules. *** Think of a highway in the states. ***What’s the highway that goes between different states called? The Interstate. •Intermolecular bonds involve the electrostatic attractive forces between molecules. •Ionic substances do not form molecules. • Therefore, intermolecular bonding only occurs in substances that form covalent bonds. (molecular Compounds) S T R O N G W E A K 3.4 Cl2.2 2.2 Types of Intermolecular Bonding Van der Waals forces can be divided into three different types London dispersion forces dipole-dipole forces hydrogen bonding WEAKEST STRONGEST London Dispersion Forces Weak attractive forces that result when the electrons of one molecule are attracted to the positive nuclei of a nearby molecule (random chance) The movement of the electrons in a molecule generates temporary positive and negative regions in the molecule A temporary dipole ***Temporary*** Let us consider the Chlorine molecule, Cl2(g): •At a particular instant, we may find that the two electrons that form the bond may be closer to one nucleus than the other. •Results in a temporary dipole with one end more negative than the other. Dipole-Dipole Forces - The positive end of one polar molecule will be attracted to the negative end of a neighbouring polar molecule. (Permanent) Hydrogen Bonding Occurs when hydrogen is bonded to a highly electronegative element (fluorine, oxygen and nitrogen) – chemistry is FON!!! The hydrogen end of the bond takes on a strong positive charge because of the exposed positive nucleus, while the other element takes on a strong negative charge This positive hydrogen will be attracted to nearby negative atoms. It appears as though the hydrogen atom bonds to different molecules. Predicting Boiling Points What affects boiling point? 1. Number of Electrons 2. Number of Carbons 3. Polarity…electronegativity difference Predicting Boiling Points In order to cause a substance or compound to boil, You must provide enough energy to break the bonds in the compound. More energy needed, higher temperature needed. In order to describe how a solid looks and behaves, you first need to determine what class of substance it is. Ionic Crystals Metallic Crystals Molecular Crystals Covalent Network Crystals Very hard (harder than ionic and molecular crystals) Brittle with very high melting and boiling points Higher than ionic and molecular Insoluble nonconductors of electricity Carbon based covalent network crystals GROUP ACTIVITY : THINGS TO WATCH FOR SATP: Standard Ambient (room) Temperature and Pressure. “Compound” refers to the joining of elements (MetalNonmetal). Ion charges of metals DO NOT have to be memorized, there on the periodic table. Ion charges of non-metals must be memorized, as they are not on the periodic table. “Molecule” refers to a molecular joining of elements (Nonmetal-Nonmetal). INDIVIDUAL----->PAIR------>GROUP---->SHARE FIRST Form pairs of two….quickly and quietly. Switch desks so that pairs are sitting next to each other. Each pair decides who is going to be A and who is going to be B. INDIVIDUAL----->PAIR------>GROUP---->SHARE A: Reads Intro and all Ionic “Stuff” (p 119-123). B:Reads Intro and all Molecular/Covalent “Stuff” (p. 119(Intro) and 124-126). As each partner reads over their pages, summarize the information (write it down). Summarize in such a way that when you are done summarizing you can explain it to your partner. First part is individual, so it should be pretty quiet. You have about 25 minutes. INDIVIDUAL----->PAIR------>GROUP---->SHARE Second Now each pair has the next 15 minutes to explain their sections to each other…BOTH must have the summarized notes. It should be a little louder now, but you should only be talking to your partners. When you are done, bring up your summaries to me and so I can have a look. INDIVIDUAL----->PAIR------>GROUP---->SHARE Third The pairs should now join up with one other pair and make a “Place Mat” of their summarized ideas. 4 People in each group. Divide into A, B, C, D, etc. WAIT! PLACE MAT ACTIVITY B’s get the markers. At the front. C’s get the Paper. At the back. A’s Divide paper into section with shared area in middle. WAIT! PLACE MAT ACTIVITY THIS TIME! Each group member summarizes the information they have in their “spot” on the place mat. You have 5 minutes WAIT! OK, NOW YOU CAN GO. Four You then rotate the “Place Mat” and check over and add (if necessary) information to the next persons summary. You will keep rotating the “Place Mat” and checking/adding until you have rotated 4 times and are back to where you started (your own). Use your time wisely. You have 15 minutes. Summarize the information your group found in the middle. Sort of a final copy.. Summarize the information your group found in the middle. Sort of a final copy..