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COLLEGE ENTRANCE TEST REVIEW CHEMISTRY CVAndaya • The Atomic Theory of MATTER Can matter be subdivided into fundamental particles? • • Democritus (460–370 BC): All matter can be divided into indivisible atomos. John Dalton: Proposed atomic theory with the following postulates: • Elements are composed of atoms. • All atoms of an element are identical. • In chemical reactions atoms are not changed into different types of atoms. Atoms are neither created nor destroyed. • Compounds are formed when atoms of elements combine. LAWS OF CHEMICAL COMBINATION • Law of constant composition: The relative kinds and numbers of atoms are constant for a given compound. • Law of conservation of mass (matter): During a chemical reaction, the total mass before reaction is equal to the total mass after reaction. • Law of multiple proportions: If two elements A and B combine to form more than one compound, then the mass of B which combines with the mass of A is a ratio of small whole numbers. e.g.: CO and CO2 (mass C = 12, O = 16, O2 = 32; ratio of O to C = 4:3, ratio of O2 to C = 8:3) • • Dalton’s theory predicted the law of multiple proportions. • A cathode ray tube (CRT) is a hollow vessel with an electrode at either end. • A high voltage is applied across the electrodes. • The voltage causes negative particles to move from the negative electrode to the positive electrode. • The path of the electrons can be altered by the presence of a magnetic field. – The amount of deflection of the cathode rays depends on the applied magnetic and electric fields. – In turn, the amount of deflection also depends on the charge to mass ratio of the electron. • In 1897, Thomson determined the charge to mass ratio of an electron to 1.76 108 C/g. Consider the following experiment: • Oil drops are sprayed above a positively charged plate containing a small hole. • As the oil drops fall through the hole, they are given a negative charge. • Gravity forces the drops downward. The applied electric field forces the drops upward. • When a drop is perfectly balanced, the weight of the drop is equal to the electrostatic force of attraction between the drop and the positive plate. • Using this experiment, Millikan determined the charge on the electron to be 1.60 10-19 C. • Knowing the charge to mass ratio, 1.76 108 C/g, Millikan calculated the mass of the electron: 9.10 10-28 g. • With more accurate measurements, we get the mass of the electron to be 9.10939 10-28 g. Consider the following experiment: • A radioactive substance is placed in a shield containing a small hole so that a beam of radiation is emitted from the hole. • The radiation is passed between two electrically charged plates and detected. • Three spots were noted on the detector: – a spot in the direction of the positive plate, – a spot which is not affected by the electric field, – a spot in the direction of the negative plate. • A high deflection towards the positive plate corresponds to radiation which is negatively charged and of low mass. This is called β-radiation (consists of electrons). • No deflection corresponds to neutral radiation. This is called gamma radiation. • Small deflection towards the negatively charged plate corresponds to high mass, positively charged radiation. This is called α-radiation. • From the separation of radiation we conclude that the atom consists of neutral, positively, and negatively charged entities. • Thomson assumed all these charged species were found in a sphere. • Rutherford carried out the following experiment: • A source of a-particles was placed at the mouth of a circular detector. • The a -particles were shot through a piece of gold foil. • Most of the a-particles went straight through the foil without deflection. • Some a-particles were deflected at high angles. • If the Thomson model of the atom was correct, then Rutherford’s result was impossible. • In order to get the majority of -particles through a piece of foil to be undeflected, the majority of the atom must consist of a low mass, diffuse negative charge - the electron. • To account for the small number of high deflections of the -particles, the center or nucleus of the atom must consist of a dense positive charge. The Nuclear Atom Rutherford modified Thomson’s model as follows: – assume the atom is spherical but the positive charge must be located at the center, with a diffuse negative charge surrounding it. • The atom consists of positive, negative, and neutral entities (protons, electrons, and neutrons). • Protons and neutrons are located in the nucleus of the atom, which is small. Most of the mass of the atom is due to the nucleus. – There can be a variable number of neutrons for the same number of protons. Isotopes have the same number of protons but different numbers of neutrons. • Electrons are located outside of the nucleus. Most of the volume of the atom is due to electrons. The Periodic Table of Elements • By 2002, there were 115 elements known. • The majority of the elements were discovered between 1735 and 1843. • How do we organize 115 different elements in a meaningful way that will allow us to make predictions about undiscovered elements? • First attempt Dmitri Mendeleev and Lothar Meyer arranged the elements in order of increasing atomic weight. Certain elements were missing from this scheme. In 1871, Mendeleev noted that As, at. no. 33, belonged properly underneath P (15) and not Si (14), which left a missing element underneath Si. He called this missing element eka-silicon. He predicted a number of properties for this element. In 1886, Ge (32) was discovered. The properties of Ge match Mendeleev’s predictions well. Atomic Structure Proton Electron Neutron Atomic Number Atomic mass Ions Isotopes Given: 56 26Fe ++ How many protons, electrons, and neutrons does the ion have? p= e= n= • Some of the groups in the periodic table are given special names. • These names indicate the similarities between group members: – Group 1A: Alkali metals. – Group 2A: Alkaline earth metals. – Group 7A: Halogens. – Group 8A: Noble gases. Periodic Table Effective Nuclear Charge: the charge experienced by an electron on a manyelectron atom. • The effective nuclear charge is not the same as the charge on the nucleus because of the effect of the inner electrons. • Electrons are attracted to the nucleus, but repelled by the electrons that screen it from the nuclear charge. • The nuclear charge experienced by an electron depends on its distance from the nucleus and the number of core electrons. • As the average number of screening electrons (S) increases, the effective nuclear charge (Zeff) decreases. • As the distance from the nucleus increases, S increases and Zeff decreases. Moseley revised Mendeleev’s PT: • The Modern periodic table: elements were arranged in order of increasing atomic number. Periodic Trends in Atomic Radii • As the principle quantum number increases (i.e., we move down a group), the distance of the outermost electron from the nucleus becomes larger. Hence, the atomic radius increases. Size of Atoms and Ions Consider a simple diatomic molecule: • The distance between the two nuclei is called the bond distance. • If the two atoms which make up the molecule are the same, then half the bond distance is called the covalent radius of the atom. • As the principal quantum number increases, the size of the orbital increases. Consider the s orbitals: • All s orbitals are spherical and increase in size as n increases. • Periodic Trends in Atomic Radii • As a consequence of the ordering in the periodic table, properties of elements vary periodically. • Atomic size varies consistently through the periodic table. • As we move down a group, the atoms become larger. • As we move across a period, atoms become smaller. • There are two factors at work: – principal quantum number, n, and – the effective nuclear charge, Zeff. • As the principal quantum number (n) increases (i.e., we move down a group), the distance of the outermost electron from the nucleus becomes larger. Hence, the atomic radius increases. • As we move across the periodic table, the number of core electrons remains constant. However, the nuclear charge increases. Therefore, there is an increased attraction between the nucleus and the outermost electrons. This attraction causes the atomic radius to decrease. Trends in the Sizes of Ions • Ion size is the distance between ions in an ionic compound. • Ion size also depends on nuclear charge, number of electrons, and orbitals that contain the valence electrons. • Cations vacate the most spatially extended orbital and are smaller than the parent atom. • Anions add electrons to the most spatially extended orbital and are larger than the parent atom. Trends in the Sizes of Ions • For ions of the same charge, ion size increases down a group. • All the members of an isoelectronic series have the same number of electrons. • As nuclear charge increases in an isoelectronic series the ions become smaller: O2- > F- > Na+ > Mg 2+ > Al 3+ Ionization Energy First ionization energy is the amount of energy required to remove an electron from a gaseous atom: • Na(g) Na+(g) + e-. The second ionization energy is the energy required to remove an electron from a gaseous ion: • Na+(g) Na 2+ (g) + e-. • The larger the ionization energy, the more difficult it is to remove the electron. • Variations in Successive Ionization Energies • There is a sharp increase in ionization energy when a core electron is removed. • Ionization energy decreases down a group. • This means that the outermost electron is more readily removed as we go down a group. • As the atom gets bigger, it becomes easier to remove an electron from the most spatially extended orbital. – Ionization energy generally increases across a period. – As we move across a period, Zeff increases. Therefore, it becomes more difficult to remove an electron. • Two exceptions: removing the first p electron and removing the fourth p electron. • The s electrons are more effective at shielding than p electrons. Therefore, forming the s2p0 becomes more favorable. • When a second electron is placed in a p orbital, the electron-electron repulsion increases. When this electron is removed, the resulting s2p3 is more stable than the starting s2p4 configuration, thus, a lower I.E. for the 4th e- in the p orbital. ELECTRON AFFINITY • Electron affinity is the energy change when a gaseous atom gains an electron to form a gaseous ion: • Cl(g) ) + e- Cl-(g) + energy or • Ar(g) + e- + energy Ar-(g) Metals and Nonmetals • Metallic character refers to the properties of metals (shiny or lustrous, malleable and ductile, oxides form basic ionic solids, and tend to form cations in aqueous solution). • Metallic character increases down a group. • Metallic character decreases across a period. • Metals have low ionization energies. • Most neutral metals are oxidized rather than reduced. • When metals are oxidized they tend to form characteristic cations. • All group 1A metals form M+ ions. • All group 2A metals form M2+ ions. • Most transition metals have variable charges. Most metal oxides are basic: • Metal oxide + water metal hydroxide • Na2O(s) + H2O(l) 2NaOH(a Nonmetals are more diverse in their behavior than metals. • When nonmetals react with metals, nonmetals tend to gain electrons: • metal + nonmetal salt • 2Al(s) + 3Br2(l) 2AlBr3(s) Alkali Metals • Alkali metal produce different oxides when reacting with O2: • 4Li(s) + O2(g) 2Li2O(s) (oxide) • 2Na(s) + O2(g) Na2O2(s) (peroxide) • K(s) + O2(g) KO2(s) (superoxide) • Alkali metals emit characteristic colors when placed in a high temperature flame. • The s electron is excited by the flame and emits energy when it returns to the ground state. Li line: 2p 2s transition Na line (589 nm): 3p 3s transition K line: 4p 4s transition • Group 6A: The Oxygen Group • Oxygen (or dioxygen, O2) is a potent oxidizing agent since the O2- ion has a noble gas configuration. • There are three oxidation states for oxygen: 2(e.g. H2O) and 1- (e.g. H2O2), ½-(e.g. KO2) • Sulfur is another important member of this group. • Most common form of sulfur is yellow S8. • Sulfur tends to form S2- in compounds (sulfides). The Halogens • Group 7A: The HalogensThe chemistry of the halogens is dominated by gaining an electron to form an anion: X2 + 2e- 2X--. • Fluorine is one of the most reactive substances known: 2F2(g) + 2H2O(l) 4HF(aq) + O2(g) H = -758.7 kJ. • All halogens are diatomic molecules, X2. • Chlorine is the most industrially useful halogen. It is produced by the electrolysis of brine (NaCl): 2NaCl(aq) + 2H2O(l) 2NaOH(aq) + H2(g) + Cl2(g). • The reaction between chorine and water produces hypochlorous acid (HOCl) which disinfects pool water: Cl2(g) + H2O(l) HCl(aq) + HOCl(aq). • Hydrogen compounds of the halogens are all strong acids with the exception of HF. As we go down Gp VIIA, the acid formed becomes stronger. CHLOROFLUOROCARBONS First CFC was synthesized in 1928 by THOMAS MIDGLEY, Jr. (who also synthesized tetraethyllead, a gasoline additive) as alternative refrigerant for NH3. It was a miracle substance. Trade name--Freon CFCs have varied uses: • Refrigerant • Foam blowing agent • Cleaning solvents for electronic parts • As metal degreaser • Sterilants for medical instruments • Propellants for pressurized chem’ls • Fire extinguisher Ozone Depleting Substances • O3 molecules are dispersed in the stratosphere, and are formed from 3 O2 + uv → 2 O3 • When brought all together to form a layer, the result amount to a thickness of 3 to 5mm. • In air, there are only 3 O3 molecules for every 10million air molecules. • Ozone is depleted where it is needed the most: 90% of all the O3 in the world is in the stratosphere. • CFCs break up into ions and atoms when they absorb UV. • CF2Cl2 + uv → CF2Cl + Cl 2Cl + 2O3 →2ClO + 2 O2 2ClO → ClOOCl ClOOCl + uv →ClOO + Cl ClOO →Cl + O2 Net eq’n: 2O3 → 3 O2 1 Cl atom could “destroy” as many as 100,000 O3 molecules Other ODS: Halons (w/ Br or F, substituting for Cl), used as fire extinguisher Methyl Bromide (more dangerous ODS), as soil fumigant Carbon Tetrachloride, a cheap, toxic solvent Methyl chloroform, a dry cleaning agent Hydrogen Chloride (from US space shuttles) Group 8A: The Noble Gases • These are all nonmetals and monatomic. • They are notoriously unreactive because they have completely filled s and p subshells. • In 1962 the first compound of the noble gases was prepared: XeF2, XeF4, and XeF6. • To date the only other noble gas compounds known are KrF2 and HArF. Chemical Bonds . • Covalent bond results from sharing electrons between the atoms. Usually found between nonmetals. • Ionic bond results from the transfer of electrons from a metal to a nonmetal. • Metallic bond: attractive force holding pure metals together. Lewis structures are used to represent covalent bonds Cl Cl H F H O H H N H H H H C H H • It is possible for more than one pair of electrons to be shared between two atoms (multiple bonds): – One shared pair of electrons = single bond (e.g. H2); – Two shared pairs of electrons = double bond (e.g. O2); – Three shared pairs of electrons = triple bond (e.g. N2). • Generally, bond distances decrease as we move from single through double to triple bonds. Multiple bonds • Multiple Bonds • It is possible for more than one pair of electrons to be shared between two atoms (multiple bonds): – One shared pair of electrons = single bond (e.g. H2); – Two shared pairs of electrons = double bond (e.g. O2); – Three shared pairs of electrons = triple bond (e.g. N2). • Generally, bond distances decrease as we move from single through double to triple bonds. • In a covalent bond, electrons are shared. • Sharing of electrons to form a covalent bond does not imply equal sharing of those electrons. • Unequal sharing of electrons results in polar bonds. • Electronegativity: The ability of one atoms in a molecule to attract electrons to itself. • Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F). • Electronegativity increases – across a period and – down a group. • Difference in electronegativity is a gauge of bond polarity: – electronegativity differences around 0 result in non-polar covalent bonds (equal or almost equal sharing of electrons); – electronegativity differences around 2 result in polar covalent bonds (unequal sharing of electrons); – electronegativity differences around 3 result in ionic bonds (transfer of electrons). Electronegativity and Bond Polarity • There is no sharp distinction between bonding types. • The positive end (or pole) in a polar bond is represented + and the negative pole -. • In general, the least electronegative element is named first. Ionic Bonding • The formation of Na+(g) and Cl-(g) from Na(g) and Cl(g) is endothermic. • NaCl(s) Na+(g) + Cl-(g) endothermic (H = +788 kJ/mol). • The formation of a crystal lattice from the ions in the gas phase is exothermic: • Na+(g) + Cl-(g) NaCl(s) H = -788 kJ/mol Electron Configurations of Ions of the Representative Elements • These are derived from the electron configuration of elements with the required number of electrons added or removed from the most accessible orbital. • Electron configurations can predict stable ion formation: – – – – – Mg: [Ne]3s2 Mg+: [Ne]3s1 not stable Mg2+: [Ne] stable Cl: [Ne]3s23p5 Cl-: [Ne]3s23p6 = [Ar] stable Covalent Bond • When two similar atoms bond, none of them wants to lose or gain an electron to form an octet. • When similar atoms bond, they share pairs of electrons to each obtain an octet. • Each pair of shared electrons constitutes one chemical bond. • Example: H + H H2 has electrons on a line connecting the two H nuclei. Ionic MgH2 Molecular Magnesium hydride FeF2 Iron(II) fluoride Mn2O3, Manganese(III) oxide H2S Hydrogen sulfide OF2 Oxygen difluoride Cl2O3 Dichlorine trioxide Resonance structures are attempts to represent a real structure that is a mix between several extreme possibilities The ozone molecule O O O O O O Resonance in Benzene • Resonance in Benzene • We write resonance structures for benzene in which there are single bonds between each pair of C atoms and the 6 additional electrons are delocalized over the entire ring: • Benzene belongs to a category of organic molecules called aromatic compounds (due to their odor). or The Chemistry of Global Warming Greenhouse Gases: those atmospheric gases which absorb infrared radiation, and help keep the earth warm. These are: H2O SO2 N2O (nitrous CO2 CH4 oxide) Global Warming Potential (GWP) CO2 = 1 N2O = 296 CH4 = 23 Abundance (ppm) CO2 = 375, N2O = 0.31 CH4 = 1.8 Anthropogenic Greenhouse Gases CF3CH2F SF6 N 2O GWP 1300 5700