Survey
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Summary of things you need to know from this first quarter. The Periodic Table - summary The modern periodic table contains the elements ordered according to their atomic number (Z), which corresponds to the number of protons in the nucleus of the atom. The general form of the periodic table is due to Mendeleev (1869) and is shown below. Hydrogen is a nonmetal The rows of the periodic table are known as periods. Moving from left to right across a period, elements have an increasing number of electrons in their outer shell (or energy level). They go from metals on the left to non-metals on the right. The columns of the periodic table are known as groups. Elements in the same group have the same number of electrons in their outer shells. Because chemical reactivity depends on the interaction of electrons in the outer shell, elements in the same group have similar chemical properties. An earlier attempt by Newlands (1863) to arrange elements by their relative atomic masses had a number of drawbacks: • The positions of some pairs of elements are reversed when ordered by mass (e.g. K and Ar) • Not all elements had been discovered at the time and Newlands left no spaces for undiscovered ones • Some ‘groups’ contained elements with obviously differing properties (e.g. the group containing oxygen and sulfur also contained iron!) Of course, at the time, scientists did not know about protons and electrons (and hence atomic number) and many elements were, as yet undiscovered. Group 1 (the alkali metals) – lithium, sodium, potassium, rubidium, caesium • are soft, reactive metals with low density which must be stored under paraffin • have atoms with one electron in the outer shell • react with water to form hydrogen gas and the metal hydroxide (which dissolves in water to give an alkaline solution), e.g. potassium + water (hydrogen oxide) _ potassium hydroxide + hydrogen • become more reactive going down the group as the outer electron (being further from the nucleus) is less strongly attracted to the positive nucleus and hence more easily lost Group 7 (the halogens) – fluorine, chlorine, bromine, iodine • have low melting and boiling points, going from fluorine and chlorine (gases), through bromine (liquid) to iodine and astatine (solids) • have atoms with seven electrons in the outer shell • are all colored, with the colors becoming deeper down the group • become less reactive going down the group as the outer electrons (being further from the nucleus) are less strongly attracted to the positive nucleus and hence additional electrons are less easily gained • will displace other halogens of lower reactivity from their salts e.g. chlorine + potassium bromide _ bromine + potassium chloride Group 0 (the noble gases) – helium, neon, argon, krypton, xenon • are unreactive gases with full outer shells of electrons Transition metals These are found in the central area of the periodic table. All other elements in the periodic table can donate or receive a fixed number of electrons and therefore make a fixed number of bonds to other elements. Transition metals can exchange a range of electrons and can therefore change the number of bonds they make. Periodic Table Vocabulary Periodic Table – a chart that organizes information about all of the known elements according to their properties. Matter – anything that has mass and volume. Element – a substance that cannot be broken down into a simpler substance by ordinary chemical means. Elemental Symbol – the single capital letter or capital letter followed by a lower case letter or letters that represents the name of the element. Atom – the smallest unit of an element that has all of the properties of the element; basic building block of matter. Molecule – two or more atoms covalently bond in fixed proportion. Compound two or more different elements bond in fixed proportion Atomic Mass – the number of protons and neutrons in the nucleus of one atom of the element. Located under the elemental symbol (at the bottom of the element box on the periodic table). Atomic Number – the number of protons contained in each nucleus of its atoms of the element. Located over the elemental symbol (at the top of the element box on the periodic table). Period – a horizontal row (left to right) in the periodic table. Group – a vertical column (up and down) on the periodic table. Reactivity – describes how likely an element is to form bonds with other elements. Valence Electrons – electrons that are located in the outermost principle energy level of an atom. Electron Cloud – the region surrounding the nucleus of an atom where the electrons are located. Ion – an atom that has gained or lost one or more electrons. Ionic Bond – a bond formed by the transfer of electrons (when valence electrons of one atom are transferred to the outer energy level of another atom). (M with non M) Covalent Bond – a bond formed when two atoms share valence electrons. (non M with non M) The Atomic Theory States: (1) All matter is composed of atoms (2) Atoms of a given element are identical (3) Atoms of different elements combine in simple whole-number ratios to form chemical compounds. Metal – an element or substance that conducts heat and electricity. Nonmetal – an element that does not conduct electricity or heat and is usually a gas at room temperature. Metalloid – an element that has some properties of a metal and some properties of a nonmetal. Inert – elements and/or compounds that when put together are unable to react chemically. The Law of Conservation of Matter – a scientific law that states that during a chemical reaction, matter cannot be created or destroyed but can be changed into a different form. Period law- The chemical properties of elements tends to repeat over specific atomic number intervals Law of definite proportions – elements within a compound have specific mass ratios Isotopes Atoms of the same element but with different numbers of neutrons 1st ionization energy The energy required to remove the most loosely held valence electron from an atom to form a positive ion. Electronegativity The tendency of an atom to withdraw electron density from a covalent bond, between it and another atom. Exothermic A process that gives out energy Endothermic A process that requires energy Periodic Trends in Atomic Properties A. Ionization Energy - the energy required to remove an electron from an atom 1. Ionization energy increases for successive electrons 2. Ionization energy tends to increase across a period a. electrons in the same quantum level do not shield as effectively as electrons in inner levels b. irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove 3. Helium has the highest First ionization energy 4. Ionization energy decreases with increasing atomic number within a group because electrons farther from the nucleus are easier to remove B. Electron Affinity / electronegativity - the energy change associated with the addition of an electron / the amount an elements pulls on electrons contained in a covalent bond 1. Affinity tends to increase across a period 2. Affinity tends to decrease as you go down in a period a. electrons farther from the nucleus experience less nuclear attraction b. Some irregularities due to repulsive forces in the relatively small p orbitals C. Atomic Radius 1 Determination of radius a. half of the distance between radii in a covalently bonded covalent atomic radii" 2r 1. Periodic Trends Radius decreases across a period because of increased effective nuclear charge due to decreased shielding Radius increases down a group because of addition of principal quantum levels The trend for increase in Atomic radius is opposite to the other two main trends IE and electron affinity (and electronegativity) An Introduction to the Periodic Table A. Organization 1. Horizontal row is called a "period" (or series) 2. Vertical column is called a "group" or "family" a. Group 1A - Alkali metals b. Group 2A - Alkaline earth metals c. Group 7A - Halogens (Gr, "salt makers") d. Group 8A - Noble gases B. Naming Elements 104 and beyond Nil = 0 un = 1 bi = 2 tri = 3 quad = 4 Pent = 5 hex = 6 sept = 7 oct = 8 enn = 9 Element 109 = un (1) nil(0) enn(9) ium = unnilennium Notes - Atoms, Molecules and Ions Fundamental Chemical Laws A. Law of Conservation of Mass 1. "Mass is neither created nor destroyed" 2. Translation: In ordinary chemical reactions, the total mass of the reactants is equal to the total mass of the products B. Law of Definite Proportion 1. "A given compound always contains the same proportions of elements by mass" 2. Translation: Compounds have an unchanging chemical formula C. Law of Multiple Proportions 1. "When two elements form a series of compounds, the ratios of the masses of the second element that combine with one gram of the first element can always be reduced to small whole numbers 2. Translation: Sometimes two elements can come together in more than one way, forming compounds with similar, though not identical formulas Dalton's Atomic Theory Please read up on Not Taught in class A. Atomic Theory 1. Each element is made up of tiny particles called atoms 2. The atoms of a given element are identical 3. Chemical compounds are formed when atoms combine with each other. A given compound always has the same relative numbers and types of atoms 4. Chemical reactions involve reorganizations of the atoms. The atoms themselves are not changed in a chemical reaction Early Experiments to Characterize the Atom A. J.J. Thomson and the Electron 1. Determined the charge to mass ratio of the electron 2. Reasoned that all atoms must contain electrons 3. Reasoned that all atoms must contain positive charges B. Robert Millikan and the Oil Drop 1. Oil drop experiments determined the charge on an electron 2. With charge information, and Thomson's charge/mass ratio, he determined the mass of an electron (9.11 x 10-31 kg) 3 . The Nuclear Atom - Rutherford's Metal Foil Experiment 1. Fired alpha particles at a sheet of thin gold foil 2. Some particles were greatly deflected a. Could not have been deflected by electrons or single protons b. Must have been deflected by a positively charged object of substantial mass, this he concluded was the nucleus. 1) Supported concept of a small, central, positive nucleus where most of the atom's mass was concentrated 2) Disproved Thomson's "plum pudding" model The Modern View of Atomic Structure: An Introduction A. Nucleus 1. Protons - positively charged 2. Neutrons - no charge 3. Small size, high density a. The mass of all of the cars in the United States in an object that would easily fit in a teaspoon B. Electrons 1. Negatively charged 2. The source of varying reactivity of different elements 3. Provide most of the atomic volume C. Atomic Number 1. Number of protons D. Mass Number 1. Number of protons + number of neutrons Relative atomic mass and Isotopes Isotopes:. Atoms with the same number of protons (same element) but different numbers of neutrons Relative atomic Mass The weighted average mass of atoms of an element Calculated by RAM = (Mass1 x % abundance1) + (Mass2 x % abundance2) + (Mass3 x %3) Abundance: The proportion of Calculated by % Abundance2 = (RAM – Mass1) x 100 (Mass2 – Mass1) Electrons a closer look Orbital Shapes and Energies A. Size of orbitals 1. Defined as the surface that contains 90% of the total electron probability 2. Orbitals of the same shape (s, for instance) grow larger as n increases B. s Orbitals 1 orbital max of 2 electrons 1. Spherical shape 2. Nodes (s orbitals of n=2 or greater) a. Internal regions of zero probability C. p Orbitals 3 orbitals max of 6 electrons 1. Two lobes each 2. Occur in levels n=2 and greater 3. Each orbital lies along an axis (2px, 2py, 2pz) D. d Orbitals 1. 2. 5 orbitals max of 10 electrons Occur in levels n=3 and greater Two fundamental shapes a. Four orbitals with four lobes each, centered in the plane indicated in the orbital label b. b. Fifth orbital is uniquely shaped - two lobes along the z axis and a belt centered in the xy plane dz2 E. f Orbitals 7 orbitals max of 14 electrons Occur in levels n=4 and greater Highly complex shapes Electron configuration A. Below Complete set of orbitals 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 Transition metal exceptions Cr 1s2 2s2 2p6 3s2 3p6 4s1 3d5 Cu 1s2 2s2 2p6 3s2 3p6 4s1 3d10 If it begins with C move one electron from s to d except for cobalt because the door is bolted. d4 and d9 are close to beginning ½ filled or filled so steal an electron from the s orbital in order to gain stability B. Write electron configuration finding the element Which period is it in ? Which block is it in? how many steps into the block is it? counting from the left Try this with Sulfur answer [Ne] 3s23P4 When writing noble gas abbreviated electron configurations, don’t forget to include all other blocks that you have to pass through when counting from the left. Quantum Numbers A. Principal Quantum Number (n) 1. Integral values: 1, 2, 3, .... 2. Indicates probable distance from the nucleus a. Higher numbers = greater distance b. Greater distance = less tightly bound = higher energy B. Angular Momentum Quantum (l) Azimuthal (this was called the "orbital quantum number" in your general chem book) 1. Integral values from 0 to n - 1 for each principal quantum number n 2. Indicates the shape of the atomic orbitals Table 7.1 Azimuthal quantum numbers and corresponding atomic orbital numbers Value of l Letter used 0 s 1 p 2 d 3 f 4 g C. Magnetic Quantum Number (ml) 1. Integral values from l to -l, including zero 2. Magnetic quantum number relates to the orientation of the orbital in space relative to the other orbitals • ml: divides the sub shell into individual orbitals which hold the electrons. Example d subshell is divided into individual orbitals as follows • -2 -1 0 +1 +2 D. Spin Quantum Number 1. Up first = + ½ Then down = - ½ Table n 1 Quantum numbers for the first four levels of orbitals in the hydrogen atom l Orbital ml designation 0 1s 0 # of orbitals 1 2 0 1 2s 2p 0 -1, 0, 1 1 3 3 0 1 2 3s 3p 3d 0 -1, 0, 1 -2, -1, 0, 1, 2 1 3 5 4 0 1 2 3 4s 4p 4d 4f 0 -1, 0, 1 -2, -1, 0, 1, 2 -3, -2, -1, 0, 1, 2, 3 1 3 5 7 Electron Spin and the Pauli Principle A. Electronic Spin Quantum Number 1. An orbital can hold only two electrons, and they must have opposite spins 2. Spin can have two values, +1/2 and -1/2 B. Pauli Exclusion Principle (Wolfgang Pauli) 1. "In a given atom no two electrons can have the same set of four quantum numbers" The Aufbau Principle and the Periodic Table A. The Aufbau Principle 1. "As protons are added one by one to the nucleus to build up elements, electrons are similarly added to these hydrogen-like orbitals B. Hund's Rule 1. "The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals Molecules and Ions A. Chemical Bonding 1. Covalent bonding - Sharing of valence electrons 2. Ionic bonding - Attraction of oppositely charged ions due to a reaction in which electrons are transferred 3. Metallic bonding –Ions in a sea of electrons B. Representing Molecules (Covalently bonded) 1. Chemical formula a. Symbols for atoms and subscripts 1) H20 2) CH4 2. Structural formula a. Bonds represented by lines Ball and Stick Space Filling C. Ions 1. Ca +ions a. Positive ions formed by the loss of electrons 2 . A n i o n s a. Negative ions formed by gaining electrons D. Ionic Bonding 1. Bond formed by the attraction between oppositely charged ions 2. Ionic bonding forms ionic solids (salts) 3. Ions can be monatomic (one atom) or polyatomic (more than one atom) E. Metallic Bonding 1. Occurs between or within metals 2 Ions in a sea of electrons 3 The valence electrons are said to be delocalized (not fixed to any one atom) Nomenclature Naming Simple Compounds A. Ionic Compounds 1. Positive ion is always named first, negative ion second Binary ionic compoundsal with non metal Contain only two elements but may have more than 2 atoms Always end in ide Magnesium chloride MgCl2 You were given a list of ions to memorize Tips for memorizing the polyatomics: 4 Find the "ate" ion (sulfate, for instance) sulfate = SO4 25 The "ite" ion always has one less oxygen than the "ate" ion sulfite = SO3 2The prefix "per" (think hyper, meaning "above") is used with the "ate" prefix to indicate one more oxygen than the "ate" ion persulfate = SO5 27 The prefix "hypo" (meaning "under" or "below") is used with the "ite" prefix to indicate one less oxygen than the "ite" ion hyposulfite = SO22Examples (Just because you can name it doesn't mean it exists!) 6 Perchlorate ClO4- Pernitrate NO4- Chlorate ClO3- Nitrate NO3- Chlorite ClO2- Nitrite NO2- hypochlorite ClO- hyponitrite NO- 2. Metals with more than one oxidation state (transition metals) must have a roman numeral to indicate the oxidation state Fe3+ = iron (III) Mn+2 = manganese (II) B. Binary Covalent Compounds 1. Must contain two elements, BOTH nonmetals a. First element 1) full element name 2) prefix only if there is more than one atom b. Second element 1) named as if it were an anion (-ide suffix) 2) always gets a prefix mono - 1 penta - 5 octa - 8 di - 2 hexa - 6 nona - 9 tri - 3 hepta - 7 deca - 10 tetra - 4 Bonding: General Concepts Types of Chemical Bonds A. Metallic Bonding Metal with Metal 1. Ions in a sea of electrons 2. Metals combined with metals (called alloys) 3. Electrons are delocalized and free to move in a conduction band B. Ionic Bonding Metal with nonmetal 1. Electrons are transferred 2. Metals react with nonmetals 3. Ions paired have lower energy (greater stability) than separated ions C. Covalent Bonds nonmetal with nonmetal 1. Electrons are shared by nuclei 2. Pure covalent (non-polar covalent) a. Electrons are shared evenly 3. Polar covalent bonds a. Electrons are shared unequally b. Atoms end up with fractional charges (1) δ+ or δElectronegativity A. Electronegativity 1. The ability of an atom in a molecule to attract shared electrons to itself B. Electronegativity Trends 1. Electronegativity generally increases across a period (why?) 2. Electronegativity generally decrease within a family (why?) C. Characterizing bonds 1. Greater electronegativity difference between two elements means less covalent character and greater ionic character Any compound that conducts an electric current when melted is an ionic compound. Bond Polarity and Dipole Moments A. Dipolar Molecules 1. Molecules with a somewhat negative end and a somewhat positive end (a dipole moment) 2. Molecules with preferential orientation in an electric field + + - - + - 3. All diatomic molecules are nonpolar Eg O2 or H2 B. Molecules with Polar Bonds but no Dipole Moment 1. Linear, radial or tetrahedral symmetry of charge distribution a. CO 2 - linear b. CCl 4 – tetrahedral c. BF 3 – Trigonal planar Guided practice Step 1 Identify the central atom Step 2 Draw a lewis dot diagram Step 3 Arranging the other atoms around the central atom Atoms should be attached clockwise in order of mass smallest first (This includes lone pairs) Step Four find lines of symmetry In reality this method only works because drawings are in 2D. In the 3D real world polar molecules are Asymmetric. Step five if the molecule has more than one line of symmetry then it is not polar Predicting Formulas of Ionic Compounds 1. Placement of elements on the periodic table suggests how many electrons are lost or gained to achieve a noble-gas configuration a. Group I loses one electron, Group II loses two, Group VI gains two, Group VII gains one.... 2. We can use the jigsaw pieces to determine formula a) Each element can be represented as a jigsaw piece b) Combine jigsaw pieces so that no unused nodes remain c) Determine how many of each element was used and construct the formula Eg Sodium (Na) with Oxygen (O) Na O Na 2O Treat polyatomic ions as single jigsaw pieces Use parenthesis to show how many polyatomic ions are required Al2(SO4)3 This formula means 2 Aluminum atoms and 3 sulfate polyatomic ions Cu(OH)2.5H2O means 1 molecule of Cu(OH)2 with 5 waters attached. The period means that the following number of molecules are attached. 3. Formulas for compounds are balanced so that the total positive ionic charge is equal to the total negative ionic charge Total positive = +6 Total negative = -6 C. Sizes of Ions 1. 2. 3. 4. Anions are larger than the parent atom Cations are smaller than the parent atom Ion size increases within a family Isoelectronic ions a. Ions with the same number of electrons b. Size decreases as the nuclear charge Z increases The trend for the ionic radii is the more negative the bigger, within the same period Lewis Structures A. Electrons and Stability 1. "the most important requirement for the formation of a stable compound is that the atoms achieve noble gas configurations 2. Duet rule a. Hydrogen, lithium, beryllium, and boron form stable molecules when they share two electrons (helium configuration) 3. Octet Rule a. Elements carbon and beyond form stable molecules when they are surrounded by eight electrons B. Writing Lewis Structures 1. Rules a. Add up the TOTAL number of valence electrons from all atoms b. Use a pair of electrons to form a bond between each pair of bound atoms. Lines instead of dots are used to indicate each pair of bonding electrons c. Arrange the remaining atoms to satisfy the duet rule for hydrogen and the octet rule for the second row elements Exceptions to the Octet Rule A. Boron Trifluoride 1. Note that boron only has six electrons around it 2. BF3 is electron deficient and acts as a Lewis acid (electron pair acceptor) 3. Boron often forms molecules that obey the octet rule B. Sulfur Hexafluoride 1. Note that sulfur has 12 electrons around it, exeeding the octet rule 2. Sulfur hexafluoride is very stable 3. SF6 fills the 3s and 3p orbitals with 8 of the valence electrons, and places the other 4 in the higher energy 3d orbital C. More about the Octet Rule 1. Second row elements C, N, O and F should always obey the octet rule 2. B and Be (second row) often have fewer then eight electrons around them, and form electron deficient, highly reactive molecules 3. Second row elements never exceed the octet rule 4. Third row and heavier elements often satisfy (or exceed) the octet rule 5. Satisfy the octet rule first. If extra electrons remain, place them on elements having available d orbitals Resonance A. Nitrate ion 1. Experiments show that all N-O bonds are equal 2. A single Lewis structure cannot represent the nitrate ion 3. A resonance structure is drawn by writing the three variant structures, connected by a double-headed arrow B. Resonance 1. When more than one valid Lewis structure can be written for a particular molecule 2. The actual structure is an average of the depicted resonance structures 3. Must have more than 2 atoms 4. Must have at least one multiple bond and more than two oxygens and or nitrogens Finding resonance structures Step 1 Draw the Lewis diagram for each of the atoms involved Step 2 Calculate the total number of electrons Step 3 Write out the molecule with the least electronegative atom in the center and connect each with one bond Step 4 For each bond drawn deduct 2 electrons from the total Step 5 Place the electrons around the outer atoms until full ( don’t forget that each bond counts as two electrons) Step 6 Move lone pair electrons from the outer atoms to make bonds with the central atom until 4 bonds are made Formal Charge Formal charge is used to find the most stable structure for a molecule. The arrangement that has the lowest range of formal charges on its atoms is the most stable. FC = Ve – (LP’s + # Bonds) FC = formal charge Ve = valence electrons Hybridization : Combination of orbitals from different subshells to produce new equal orbitals Number of electron directions equal number of combined orbitals Number of different directions of e- pairs can be found must match number of orbitals merged LP = lone pair electrons The electrons are in 4 different directions so the nitrogen needed 1 s combined with 3p’s to make 4 new sp3 hybridized orbitals Molecular Structure: The VSEPR Model A. Valence Shell Electron Pair Repulsion (VSEPR) 1. The structure around a given atom is determined principally by minimizing electron-pair repulsions 2. Non-bonding and bonding electron pairs will be as far apart as possible Arrangement of Electron Pairs Around an Atom Yielding Minimum Repulsion # of Electron Shape Example Lewis dot diagram Pairs 2 Linear Bond angle 180o 3 Trigonal Planar Bond angle 120o 4 Tetrahedral Bond angle 109.5o 5 Trigonal bipyramidal Bond angle 120o Bond angle 180o 6 Octahedral Bond angle 90o B. Effect of Unshared Electron Pairs Complete 1. The ideal tetrahedral angle is 109.5° Comparison of Tetrahedral Bond Angles Compound Structure Angle between Hydrogens Methane 4 0 Lone 109.5° Bonding pairs electron pairs Ammonia 3 1 lone bonding pair pairs 107° Water 2 2 lone bonding pairs pairs 104.5° 2. Lone (unshared) electron pairs require more room than bonding pairs (they have greater repulsive forces) and tend to compress the angles between bonding pairs 3. Lone pairs do not cause distortion when bond angles are 120° or greater 4. Substitution of bonding pairs with lone pairs decreases the bond angle by 2.5o C. VSEPR and Multiple Bonds 1. For the VSEPR model, multiple bonds count as one effective electron pair 2. When a molecule exhibits resonance, ANY of the resonance structures can be used to predict the molecular structure using the VSEPR model D. Molecules Containing No Single Central Atom 1. Apply the principal of distancing shared and unshared electron pairs 2. Look at real 3-dimensional, rotatable models to develop predictive skills