Honors Chem - Bonding Unit PDF

Honors Chem - Bonding Unit A chemical bond is an electrostatic attraction: -the protons of one atom are attracted to the e- of a neighboring atom. -a mutual attraction between the nuclei and valence e- of different atoms that binds the atoms together The...

Honors Chem - Bonding Unit A chemical bond is an electrostatic attraction: -the protons of one atom are attracted to the e- of a neighboring atom. -a mutual attraction between the nuclei and valence e- of different atoms that binds the atoms together The electrons are THE subatomic particle involved in bonding. Reminders: -subatomic particles = protons, neutrons, electrons -protons and neutrons are found inside the nucleus -electrons are located outside the nucleus. an orbital is the region where electrons are most likely to be found. -the valence shell is the outermost energy level -the valence electrons are the electrons located in the valence shell -8 is the maximum number of valence electrons -atoms with the same # of valence e- behave similarly -atoms with the same # of valence e- will bond/react similarly -atoms found in the same group on the periodic table have the same # of valence e- and therefore will bond/react similarly. Common questions with their answers: QUESTION ANSWER How can we tell, in terms of atomic structure, which Atoms of elements with the same # of valence elements/atoms will bond in similar ways? electrons will bond in similar ways. How can we tell, in terms of the periodic table, Atoms of elements located in the same group will which elements/atoms will bond in similar ways? bond in similar ways. Why do atoms bond? Individual atoms that do not have full valence shells have a high amount of PE, but are not stable. Nature favors low PE (high stability). So, atoms bond to become more stable. When is an atom stable? An atom is stable when it has a full valence shell. Atoms of which elements are stable (and therefore, Noble gases/inert gases/group 18 do not readily react or bond)? Know these three types of bonds: ionic, covalent, metallic When a question asks “ what type of bonding is present in…” your answer will need to be one of the following: ionic, covalent, metallic, or ionic & covalent (we’ll discuss later). When atoms bond, their valence e- are redistributed in different ways to make the atoms more stable. Bond type is determined by the way in which the valence e- are redistributed.  Ionic bond = e- transferred Covalent bond = e- shared Metallic bond = "sea of freely moving (mobile) e-" Octet rule: atoms gain, lose, or share e- to acquire a stable octet (8 valence e-) Exceptions to the octet rule: Helium is a noble gas. Helium has a full valence shell. Helium is stable even though it does not have 8 valence e-. Stability and potential energy: Bonds have/store PE (stored energy) Individual atoms have more PE (less stability) than compounds. Atoms want to become part of compounds so that the atoms will have less PE (more stability). Atoms Compounds high PE low PE low stability high stability Q: Explain what happens, in terms of energy, when a bond is formed and when a bond is broken. A: When a bond is Formed, energy is Released. Energy must be Absorbed to Break a bond. Break Absorb Release Form (Perhaps chemistry sometimes makes you want to BARF!) I2 --> I + I shows breaking a bond (BArf) I + I --> I2 shows forming a bond (baRF) energy + I2 --> I + I I + I --> I2 + energy A bond will be formed when atoms combine (or ions of opposite charges attract). BA/RF: Break (bond or attraction) Absorb (energy) energy will be on the left of the arrow Release (energy) Form (bond or attraction) energy will be on the right of the arrow 2 Reminders: -Electronegativity is a measure of an atom's ability to attract electrons. -Fluorine is the most electronegative element (has the highest electronegativity). Fluorine has an electronegativity of 4.0 -Fluorine is the most reactive nonmetal. -Compounds are atoms of 2 or more different elements chemically combined in a fixed proportion. -Binary compounds are compounds that are made up of only 2 elements (this has nothing to do with the number of atoms/ions of each of these elements). -Ternary compounds are compounds that are made up of 3 or more different elements (again this has nothing to do with the number of each). -COMPOUNDS ARE NEUTRAL! -Ions are charged particles cations are positively charged ions anions are negatively charged ions Ionic Bonding -Ionic bonds are formed when e- are gained or lost (transferred). -A bond may be ionic when its electronegativity difference is ≥1.7 (a bond may still be ionic even if it does not fall into this GUIDE). -Ionic bonding exists in ionic compounds. (Ionic bonds form ionic compounds) Q: What is an ionic compound? A: a compound that is composed of positive and negative ions that are combined so that the # of positive and negative charges are equal b/c COMPOUNDS ARE NEUTRAL! (Ions are formed and attract one another) Key phrases for ionic: "Transfer of Electrons" "Salts" Crystal Lattice structures Metals loses electron(s) and Nonmetal gains electron(s) Number of electrons gain must equal the number of electrons lost Compound: 2 or more elements chemically combined in a specific or definite ratio. -This definite ratio is based on the octet rule, and the fact that the number of electrons lost must equal the number of electrons gained! 3  Dot diagrams showing ratio of ions: Groups 1 and 17 Groups 2 and 17 Example: Na and Br Example: Mg and Cl Ratio of ions (positive to negative): Ratio of ions (positive to negative): Chemical formula: Chemical formula: Name: Name: LDD: LDD: Groups 1 and 16 Groups 2 and 16 Example: K and O Example: Ca and S Ratio of ions (positive to negative): Ratio of ions (positive to negative): Chemical formula: Chemical formula: Name: Name: LDD: LDD: Groups 1 and 15 Groups 2 and 15 Example: Rb and P Example: Sr and N Ratio of ions (positive to negative): Ratio of ions (positive to negative): Chemical formula: Chemical formula: Name: Name: LDD: LDD: Groups 13 and 17 Groups 13 and 15 Example: Al and F Example: Al and N Ratio of ions (positive to negative): Ratio of ions (positive to negative): Chemical formula: Chemical formula: Name: Name: LDD: LDD: Groups 13 and 16 Example: Al and O Ratio of ions (positive to negative): Chemical formula: Name: LDD: 4 Q: How can you identify a binary ionic compound? A: A binary ionic compound is made up of only 2 elements: one metal and one nonmetal. example: Explain, in terms of classification of elements, why MgBr2 is an ionic compound. (answer: Mg is a metal and Br is a nonmetal) Q: What is a salt? A: any ionic compound in which H+ is not the positive ion and OH- is not the negative ion (not just table salt)! Monatomic Ions: Monatomic ions are ions that are formed when an individual atom has either gained or lost one or more electron(s). Metals: -have a relatively low number of val e- -lose their val e- to become + ions -M atoms have larger radii than their corresponding ions examples: sodium ion: calcium ion: aluminum ion: -ions of metals that have more than one possible oxidation state must get a Roman numeral as part of the ion name to identify the ion. The Roman Numeral IS the charge (oxidation number) of the ion. examples: Cu+1 = copper (I) Cu+2 = copper (II) Sn+2 = tin (II) Sn+4 = tin (IV) +2 Is Zn called the zinc ion or the zinc (II) ion? Why? Nonmetals: -have a relatively high number of val e- -gain e- to form negative ions -NM atoms have smaller radii than their corresponding ions. -monatomic NM ions end in -ide. -charge of the ion is the first oxidation number listed in the box for the nonmetal examples: iodide ion: phosphide ion: oxide ion: 5 Polyatomic Ions Polyatomic ions are charged groups of covalently bonded atoms -ions that are made up of 2 or more atoms joined together by covalent bonds and have a charge. -memorize all of your ions from the start of the year! -the charge of a polyatomic ion applies to the entire group -- not just the rightmost element. S2O32- does not mean that the O has a charge of 2-. The 2- applies to the entire group of atoms (S2O3)2- CrO42- = (CrO4)2- PO43- = (PO4)3- OH- = (OH)- -The names of polyatomic ions never change. Keep the name that is on table E. -When there is no number next to the charge, then the number is really 1. So, ternary ionic compounds will always have a polyatomic ion. -This could be one polyatomic ion and one monatomic ion. Examples: -This could be two polyatomic ions. Examples: ***NH4+ (ammonium) is an “honorary metal”. It is the only positive polyatomic ions you will see in compounds that are made of NM only. 6 Understanding Subscripts for Ionic Compounds -No subscript = 1 -Subscripts are written to the right of the element symbol. -For ionic compounds, the subscript tells you the number of each type of ion. Binary Examples: NaF means you have 1 sodium ion (Na+) and 1 fluoride ion (F-). MgCl2 = you have 1 magnesium ion (Mg2+) and 2 chloride ions (Cl-). Cu2O = you have 2 copper (I) ions (Cu+) and 1 oxide ion (O2-). Notice that when you add all the charges up in an ionic compound the total equals zero b/c compounds are neutral! NaF: (+1) + (-1) = 0 MgCl2: (+2) + (-1) + (-1) = 0 Cu2O: (+1) + (+1) + (-2) =0 Ternary Examples: KMnO4 means you have 1 potassium ion (K+) and 1 permanganate ion (MnO4-) Ca(OH)2 = means you have 1 calcium ion (Ca2+) and 2 hydroxide ions (OH-) Na3PO4 = means you have 3 sodium ions (Na+) and 1 phosphate ion (PO43-) Al2(Cr2O7)3 = means you have 2 aluminum ions (Al+) and 3 dichromate ions (Cr2O72-) Notice that when you add all the charges up in an ionic compound the total equals zero b/c compounds are neutral! KMnO4: (+1) + (-1) =0 Ca(OH)2: (+2) + 2(-1) =0 Na3PO4: 3(+1) + (-3) =0 Al2(Cr2O7)3: 2(+3) + 3(-2) =0 Lewis Dot Diagrams for Ionic Compounds -simply draw the Lewis Dot Diagrams for each ion. Examples: NaF (sodium fluoride) Al2O3 (aluminum oxide) MgCl2 (magnesium chloride) 7 Ionic bonds are not the same thing as ionic compounds. An ionic bond is the attraction between 2 ions. Ionic compounds will either contain ionic bonds only OR ionic and covalent bonds: IONIC BONDS ONLY IONIC AND COVALENT BONDS The attraction between a positive The attraction between the positive ion and negative ion is monatomic ion and a negative ionic. The attractions within the polyatomic ion(s) are monatomic ion is ionic covalent. These will always have at least one polyatomic ion Al2O3 MgCl2 KI CaCO3 NaOH CuSO4 Properties of Ionic Substances (ionic compounds): -Hard Q: What makes something a good conductor?/What allows for conductivity? A: Something will conduct when it has freely moving charged particles. -Good conductors in the liquid phase(of heat and electricity) b/c the ions are free to move. -Good conductors in solutions b/c the ions are free to move. Reminder: (aq) = aqueous solution = homogeneous mixture = mixed with water -Nonconductors in the solid phase b/c the ions are not free to move. The ions are stuck in the crystal lattice Does NOT conduct as a solid but does conduct as (l) or (aq) will always indicate an IONIC COMPOUND! -high melting and boiling points -dissolve in water (and other polar substances) The greater the difference in electronegativity, the “more Ionic” or “more stable” the bond is. 8 Naming/Formula Writing for Ionic Compounds = Nomenclature Use the stock system 1. Metal (or ammonium) comes first (positive ion) 2. Non Metal comes last (negative ion) 3. Binary compounds end in IDE (compounds made up of only 2 elements) 4. Ternary compounds end in ITE or ATE usually (compounds made up of 3 or more elements; have polyatomic ions) The only negative polyatomic ions that end in -ide are hydroxide, peroxide, and cyanide. All other “ide” means the ion is a monatomic nonmetal ion. 5. Sum of oxidation states must always equal zero (number of electrons lost must equal number of electrons gained) 6. If you need more than one of the same polyatomic ion to satisfy rule #5 you must put the polyatomic ion in parenthesis. 7. If the metal or the element acting as the metal has more than one possible oxidation state (usually the transition) you MUST use roman numerals to identify the oxidation state. *** 8. Ionic compounds must be in empirical form (most reduced ratio) *** Memorize the oxidation states for Metals with only 1 Oxidation State Fe = +2 or +3 Co = +2 or +3 Therefore they never need Roman Numerals! Ni = +2 or +3 Mn= +2 or +3 (and more) Zn = +2 Al+3 Cu = +1 or +2 Cr=+2 or +3 Ag = +1 Au = +1 or +3 Cd = +2 Sn = +2 or +4 W = +6 Pb = +2 or +4 Old System for naming and formula writing for Ionic Compounds: This is only for recognition. You will not name compounds this way. -based on the Latin name for the metal, with the suffix changed -the suffix will indicate what the oxidation state for the metal -you will never name compounds using this system…this is only to recognize if the old name is given Suffix ____________ic == Higher of the oxidation state Examples: ____________ous == Lower of the oxidation state Cupric = Cu+2 Ferric = Fe+3 Cuprous = Cu+1 Ferrous= Fe +2 Name the following: 1. MgCl2 2. NaF 3. Ca(OH)2 4. FeO 5. Fe2O3 6. AuClO3 7. AgClO3 9 Covalent Bonding Reminders: -3 types of bonds: ionic, covalent, metallic -bond type is determined by how e- are redistributed: ionic (transfer) covalent (share) metallic (sea of mobile e-) -why bond? to become stable -BARF (break absorb release form) -gp 18 gen. don't bond -Ionic -- transfer e- -- M and NM _______________________________________ Covalent = share e- = NMs only Q: Between what types of atoms will a covalent bond form? A: nonmetals Q: How are electrons redistributed in covalent bonds? A: e- are shared Q: In what ways can electrons be shared? A: equally or unequally Q: What do we call a covalent bond in which electrons are shared equally? A: nonpolar Q: What do we call a covalent bond in which electrons are not shared equally? A: polar Q: How can you remember this? A: Q: What makes up one covalent bond? A: 1 covalent bond = 2 electrons being shared = 1 pair of e- being shared One covalent bond is between 2 NM atoms Q: How can we tell if the two atoms of a covalent bond will share electrons equally or unequally? A: compare the ability to attract e- of the 2 different atoms. Q: What measures an atom's ability to attract electrons? A: electronegativity 10  Reminders: Electronegativity is: -the ability to attract electrons -the affinity for electrons -the love of electrons -the likeliness to gain electrons -Fluorine has the highest electronegativity = 4.0 Q: How can we determine an atom's electronegativity? A: look it up on Table S (Regents) and know your trends! Q: Based on electronegativity values, how can we tell if two atoms will share electrons equally or unequally? A: -when the 2 atoms have the same electronegativity value the electrons will be shared equally. -when the 2 atoms have different electronegativity values the electrons will be shared ≠. It might help you to think of electronegativity as how strongly an atom is pulling on the shared pair of electrons in a covalent bond. Q: Based on the electronegativity difference GUIDE, what electronegativity difference will a bond be considered nonpolar (share e- =)? A: -a bond is nonpolar when it exists between two atoms of the same electronegativity (same NM) Q: Based on the electronegativity difference GUIDE, what electronegativity difference will a bond be considered polar (share e- ≠)? A: Bond polarity is based off electronegativity difference! The bigger the difference in electronegativity, the more polar the bond. The smaller the difference in electronegativity, the less polar the bond. 11 Nonpolar Covalent Bonds -share e- equally (same electronegativity) -same NM -H2 O2 F2 Br2 I2 N2 Cl2 Q: What do F2, Cl2, Br2, and I2 have in common? A: Draw the LDD (e- dot diagram) for 2 atoms of fluorine (F). Draw the LDD for a fluorine molecule (F2). The 2 dots in the middle of the 2 symbols represent e- being shared by both atoms. Structural formulas do not show any dots. A structural formula shows the arrangement of atoms in a molecule by using a dash to indicate each covalent bond. Remember, 1 covalent bond is really made up of 2 e- (1 pair of e-). Draw the structural formula for a molecule of fluorine. Draw the LDD/e- dot diag and the structural formula for a molecule of chlorine (Cl 2). Draw the LDD and structural formula for a bromine molecule (Br2). Draw the LDD and structural formula for an iodine molecule (I2). 12 Draw the LDD and structural formula for a molecule of hydrogen (H2). Remember that Hydrogen needs 2 valence e- to be full and NOT 8 b/c hydrogen only has one shell. ------------------------------- Draw the LDD and structural formula for molecular oxygen (O2). The oxygen molecule has a double bond. The two atoms share 4 e- (not just 2). ----------------------------------- Draw the LDD and structural formula for a molecule of nitrogen (N 2). The nitrogen molecule has a triple bond. The two atoms share 6 e- (not just 2). Be careful when you are asked to draw LDD for these diatomic elements. If you are asked to draw an ATOM you draw only 1 atom with its correct number of valence e-. If you are asked to draw a MOLECULE you draw the diatomic molecules with their bond(s) as we did above. Memorize the LDD and structural formulas for H2, O2, F2, Br2, I2, N2, Cl2 and that those molecules have nonpolar bonds. Single bond Double bond Triple bond Strength (bond Weakest strongest energy) Bond length Longest shortest The stronger the bond, the more energy is released when it is formed and the more energy is absorbed when it is broken. 13 Reminders up to this point: Covalent: NM and NM Share electrons 1 covalent bond = 2 e- shared = 1 pair of e- Nonpolar bond: e- shared equally Polar bond: e- not shared equally BARF Nonpolar bond=same element=same electronegativity so the electronegativity difference = 0 Polar Covalent Bonds -share e- unequally -different nonmetals therefore different electronegativities One covalent bond is between 2 NM atoms. One molecule may have many, many, many bonds! Q: Is the bond between H and F polar or nonpolar? Why? A: polar b/c the elements are different NM with different electronegativity values. Fluorine pulls the shared pair of electrons much closer than hydrogen does. F "pulls with a force of" 4.0 while H "pulls with a force of" 2.2 Draw the LDD and structural formula for HF. Draw the LDD and structural formula for HCl. Draw the LDD and structural formula for HBr. Draw the LDD and structural formula for HI. 14 LDD for polyatomic ions: remember that the electrons are shared between the atoms of the polyatomic ion. The species formed is an ion (has a charge) and forms an ionic bond with other ions. Example: Hydroxide ion Q: What type of bonding is present in MgSO4? A: both ionic and covalent b/c the atoms of polyatomic ions are held together with covalent bonds, but the polyatomic ions bond with other ions via ionic bonds. Q: What is a coordinate covalent bond? A: A covalent bond in which both electrons in the bond pair have been contributed by the same atom. In this bond, one of the atoms comes with no valence electrons to share. That atom is a “free-loader”. The “free-loader” atom is usually a hydrogen ion (H+). A molecule with a lone pair of electrons can theoretically form a coordinate covalent bond. This is different from a molecule or ion that already has a coordinate covalent bond. Note the difference. H2O and NH3 can form coordinate covalent bonds. H3O+ and NH4+ each have a coordinate covalent bonds. These are not the only two examples. 15 LDD for Covalent compounds/ Molecules, and polyatomic ions Reminder: The outermost s and p sublevel electrons are the valence electrons! Exceptions to the octet rule There are many exceptions to the octet rule! 1) H only needs 2 electrons to be complete (He only needs 2 electrons since it only has one energy level). 2) Beryllium (Be) will form covalent bonds NOT ionic bonds with halogens Be will form 2 covalent bonds with a total of 4 electrons (not 8) 3) Boron (B) may form 3 covalent bonds with a total of 6 electrons (not 8). BF3 4) Atoms can form expanded octets to stable shells of 10, 12, or 14. P, S, Cl, Br, and I all do this. PCl5 (P has 5 bonds totaling 10 electrons) SF6 (S has 6 bonds totaling 12 electrons) Drawing LDD for molecules and polyatomic ions (each have covalent bonds between atoms) Find the central atom if possible. LDD only show 2d. Molecules are 3d. VSEPR model (theory) predicts 3D shapes: valence shell electron pair repulsion The overall shape of a molecule is determined by its bond angles. 1. Draw ldd LDD Count-up method 1. Count up total # of val e- involved in bonding (don't forget to add or subtract e- for polyatomic ions as needed) 2. Connect atoms to central atom with single bonds (FOR NOW) 3. Fill in lone pair e- to get the total found in step one. 4. Move lone pairs into multiple bond position(s) if needed. (C,O,N,S are the most common atoms to form multiple bonds) 2. Determine the total number of e- domains -total number of e-domains tells you the e domain geometry -e domains want to be as far apart as possible 4 possible e- domains: -, =, ≡, and **. Each of these is its own “1” domain. e- domain geometry ≠ molecule geometry 3. From the total # of e- domains, determine the # of bonding domains and the # of nonbonding (lone pair) domains. 16  MEMORIZE: # bonded and # lone Shape and bond angle 2 atoms only (H2, O2, HF etc.) Linear 180° 2 bonded and 0 lone pairs Linear 180° 3 bonded and 0 lone pairs Trigonal planar 120° (need to redraw) 2 bonded and 1 lone pair Bent 120° (slightly less than 120°) (need to redraw) 4 bonded and 0 lone pairs Tetrahedral 109° 3 bonded and 1 lone pair Trigon pyramidal 107° (slightly less than 109°) 2 bonded and 2 lone pairs Bent 105° (need to redraw) (slightly less than 107°) 5 bonded and 0 lone pairs Trigonal bipyramidal 90°120° and 180° 6 bonded and 0 lone pairs Octahedral 90° and 180° 17 1. H2O 2. NH3 3. CH4 4. F2 5. PCl3 6. **CO2** 7. OH- 8. ClO2- 18 9. NH2- 10. HCN 11. BF3 (B is an exception to the octet rule! B needs 6 electrons to be stable). 12. PCl5 (P is an exception to the octet rule! In this case P needs 10 electrons to be stable). 13. BeCl2 (Be is an exception to the octet rule!) 19 Resonance structures -resonance occurs when a molecule or polyatomic ion has more than one way to draw the ldd. -resonance often occurs when a molecule or polyatomic ion has multiple bonds -use between resonance structures -the "real" existence of a molecule or ion that shows resonance is a picture that can't be represented by one picture. 1. O3 2. SO2 3. SO3 20 Recap: Linear e- domain geometry (2 areas of e-) CO2 BeCl2 Trigonal planar e- domain geometry (3 domains = 3 areas of e-) Trigonal planar molecule geometry bent molecule geometry BF3 (3 bonding domains) SO2 (2 bonding domains and 1 nonbonding domain) Tetrahedral e- domain geometry (4 domains = 4 areas of e-) Tetrahedral trigonal pyramidal bent Molecule geometry molecule geometry molecule geometry CH4 NH3 H2O (4 bonding domains) (3 bonding, 1 nonbonding domain) (2 bonding, 2 nonbond) Trigonal bipyramidal e- domain Octahedral e- domain geometry Draw the carbonate ion: 21 Why do orbitals form these shapes? Orbital overlap (good for easy diatomic molecules) F2 1s22s22p5 1s22s22p5 Hybrid orbitals/hybridization -atomic orbitals mix to form new orbitals called hybrid orbitals 1. sp hybrid orbital BeF2 2. sp2 hybrid orbital BF3 3. sp3 hybrid CH4 4. sp3d 5. sp3d2 Sigma bonds Pi bonds Single bonds The 2nd bond in a double bond The first bond in a double or triple bond The 2nd and 3rd bonds in a triple bond Overlapping orbitals 22 Determining the polarity of a molecule Molecule polarity is determined by molecular shape and by the polarity of the bonds dipole = polar bond AND dipole=polar molecule H-Cl If a molecule has more than 2 atoms: if the dipoles cancel, the molecule is nonpolar if the dipoles do NOT cancel, the molecule is polar Q: What is a molecule? A: the smallest particle capable of independent motion. Noble gases are molecules that are made up of only one atom. All other molecules are made up of two or more atoms covalently bonded together. Q: Are molecules elements or compounds? A: Molecules may be elements or compounds. Q: What are compounds that have covalent bonding called? A: molecular or network Reminders: -1 covalent bond is between 2 atoms -1 molecule may have many, many, many bonds Molecules have polarity Bond polarity is completely different from molecule polarity Q: What determines a molecule's polarity (polar or nonpolar)? A: - the distribution of charge -whether the e- are distributed evenly (equally) or unevenly (unequally) -whether the molecule has a symmetrical or asymmetrical distribution of charge Q: In terms of symmetry and distribution of charge, when will a molecule be considered a polar molecule? A: asymmetrical and uneven distribution of charge Q: In terms of symmetry and distribution of charge, when will a molecule be considered a nonpolar molecule? A: symmetrical and even distribution of charge SNAP Symmetrical Nonpolar Asymmetrical Polar 23  Drawing the LDD for a molecule will help you determine its polarity. Draw the LDD and structural formula for a molecule of H2O. Q: Does water have polar or nonpolar bonds? Explain your answer. A: polar bonds. Each bond between H and O is polar because H and O have different electronegativity values, and therefore do not pull equally on the shared pair of e-. Q: Is water a polar or nonpolar molecule? Explain your answer. A: polar molecule. There is an uneven and asymmetrical distribution of charge. When drawing LDD for molecules, start by drawing the LDD for the individual atoms. For each of the following molecules draw the LDD and structural formula. State if the molecule has polar or nonpolar bonds. State if the molecule is polar or nonpolar and why. 1. Ammonia (NH3) 2. Methane (CH4) 3. Phosphorus trichloride (PCl3) 4. Carbon dioxide (CO2) Memorize the LDD, structural formula, and polarity of the molecules above. Molecule polarity is based on “distribution of charge”. A molecule is polar because there is an asymmetrical distribution of charge. A molecule is nonpolar because there is a symmetrical distribution of charge. 24 Naming and Formula Writing for Molecules Nomenclature Based in using Prefixes to indicate the number of atoms in the compounds First element: -gets its full name -only gets a prefix if the prefix is not “mono” Second element: -always gets a prefix (the textbook will saw the second element doesn’t get a mono, but go by the notes) -always ends in -ide Memorize the prefixes! Prefixes: 1-Mono 6-Hex 2-Di 7-Hept 3-Tri 8-Oct 4-Tetra 9-Non 5-Pent 10-Dec Examples: CO CO2 NO2 N2O5 Sulfur Trioxide Dihydrogen monoxide 25 Intermolecular forces (again) (This lesson consists mostly of reminders) Q: What are intermolecular forces? A: The forces of attraction that exist between neighboring molecules. (Bonds exist within a molecule) There are different types of intermolecular forces each with its own relative strength. 1. Dispersion Forces (London Dispersion Forces) -weak intermolecular forces -these forces keep nonpolar molecules together -these forces allow nonpolar molecules to exist in states other than gas at low temperatures. 2. Dipole-Dipole Attraction Q: What is a dipole? A: A polar molecule A polar molecule has 2 poles. -one pole (end) is slightly negative and the other end is slightly positive. Q: What determines which end of a polar molecule is slightly negative and which pole (end) is slightly positive? A: the distribution of charge. Q: In terms of electronegativity, which pole will be slightly negative? A: The pole that has a higher electronegativity (pulls more on the shared pair of e-) Q: In terms of electronegativity, which pole will be slightly positive? A: The pole that has a lower electronegativity (pulls less on the shared pair of e-) -Dipole-dipole forces are stronger than dispersion forces. -These intermolecular forces keep polar molecules together. -This attraction occurs when the positive end of one dipole (polar molecule) is attracted to the negative end of a neighboring dipole. 3. Hydrogen Bond (Hydrogen bonding) -the name "hydrogen bond" is a misnomer. Even though the word bond is in the name, this is NOT a bond. Hydrogen bonds are intermolecular forces. -This is a super strong dipole-dipole attraction that gets its own name. -This bonding is found in H2O, NH3, and HF. It is the attraction between the H of one molecule to F, O, or N of a neighboring molecule. The stronger the IMFs the higher the boiling point! 26 4. Molecule-Ion attraction -not a true "intermolecular" force b/c it is between a polar molecule and an ion (not between 2 molecules). Reminder: An ion is a charged particle -A polar molecule (ex = H2O) is attracted to an ion. The positive end of the polar molecule is attracted to the negative ion. The negative end of the polar molecule is attracted to the positive ion. Which end of H2O is slightly positive? H Which end of H2O is slightly negative? O https://www.youtube.com/watch?v=9YwdeEDrfPI What Are Intermolecular Forces | Properties of Matter | Chemistry | FuseSchool 27 Metallic Bonding -Metal atoms only -"sea of freely moving/mobile e-" -cations immersed in a sea of electrons Properties 1. Conductivity Q: What allows for a substance to conduct? A: In order for something to conduct it must have freely moving charged particles (charged particles will either be ions or electrons depending on the sample). Q: Will substances with metallic bonding conduct? Why or why not? A: Metals -excellent conductors in any phase b/c metals have freely moving electrons Q: Will substance with ionic bonding conduct? Why or why not? A: Ionic (salt) a) solid - does NOT conduct because the ions are not free to move. The ions are stuck in the crystal lattice (fixed rigid geometric pattern). b) liquid - DOES conduct because the crystal lattice has been destroyed. The ionic liquid conducts b/c it has freely moving ions (the ions are free to move). c) aqueous (aq) - DOES conduct b/c the ions are free to move. (Reminder: (aq) = mixed with water) Q: Will substances with covalent bonding conduct? Why or why not? A: Covalent -does NOT conduct b/c these substances do not have freely moving charged particles. Covalent substances do not have freely moving ions or freely moving electrons. 28 2. Melting/Boiling Points Q: How do boiling and melting points relate to the strength of attractive forces between particles? A: (Red Rover) The stronger the forces of attraction between neighboring particles the higher the melting and boiling points will be b/c more energy will be needed to break those forces of attraction. (weaker forces = lower mp/bp) Q: How do the relative mp/bp compare in substances with different types of bonding? A: Metallic - very high Ionic - very high Network solid - very high. Think diamond! Molecular - low -w/ hydrogen bonding = relatively high mp/bp in comparison with other covalent substances -w/ dipole-dipole attraction = not as high as those with hydrogen bonding, but still higher than those with dispersion forces -w/ dispersion forces = very low Relating MP/BP, strength of attractive forces, mass and nearness to ideal: Within nonpolar molecules: Lower mass = smaller momentary dipole=weaker IMF=faster to effuse=closer to ideal H2 (2g/mole) 20.K N2 (28g/mole) 77K O2 (32g/mole) 90.K 3. Solids Metallic - very hard Ionic - hard Molecular - soft Network solids: Diamonds (C)-very hard Graphite (C) Sand = Silcon DiOxide (SiO2) Silicon Carbide (SiC) SPLash into H2O. H2O is molecular. Properties of molecular include soft, poor (conductors), and low (mp/bp). MoSoPoLo (MOlecular SOft POor conductors LOw mp/bp) LIKE DISSOLVES LIKE Nonpolar substances will dissolve in other nonpolar substances. Polar substances will dissolve in other polar substances. Ionic dissolves in polar due to the molecule-ion attractions. 29 Substances and their properties: Metals Ionic Substances Network solids Molecules (metallic (Ionic (molecular substances) Compounds) substances) Solids are hard hard Hard (graphite is soft an exception) Boiling point Very high Very high Very high Low Conductivity Always great (s) DO NOT DO NOT DO NOT conductors in all (l) and (aq) DO (graphite is an phases due to Based off of if exception) mobile e- ions are mobile Type of Bonding Metallic bonds Ionic Bonds Covalent Covalent OR Ionic and Covalent https://www.youtube.com/watch?v=oNBzyM6TcK8&list=PL65159266CFC74682&index=16 Mark Rosengarten’s “What Kind of Bonds are These?” https://www.youtube.com/watch?v=cgiNk94XyaI&list=PL65159266CFC74682&index=17 Mark Rosengarten’s “Hydrogen Bonds” 30