Chem Reviewer PDF - 1.1 What is Chemistry

chem reviewer 1.1 What is chemistry? Study of transformation and behavior of matter. What is matter? Occupies space, has mass, and consists of atoms and molecules. Macroscopic - can be seen by the naked eye Mi...

chem reviewer 1.1 What is chemistry? Study of transformation and behavior of matter. What is matter? Occupies space, has mass, and consists of atoms and molecules. Macroscopic - can be seen by the naked eye Microscopic - can’t easily be seen by the naked eye, need a microscope. Atomic - nanoscale or atomic scale, atoms and molecules can’t be seen with the naked eye. Purpose of chemistry: design thoughtful experiments and make careful observations of macroscopic matter. careful observation is required. all ideas are open to challenge. Scientific method: examining results and seeing if there is another way to interpret the data. Qualitative Observations - observations with no numerical measurement. Ex: color, size (big, small), temperature (hot, cold), smell Quantitative observations - observations with numerical measurement. measurements are expressed with units. units are characterized by dimensions. Ex of dimensions: temperature, mass, length, volume, time, distance, density Ex of unit: meter, celsius, candela, seconds chem reviewer 1  1.2 Classification of matter: Element - a simple type of matter, a pure substance that cannot be broken down further, we have 118 elements Atom - smallest, indivisible unit of an element Chemical Compound - a combination of two or more elements that are chemically bonded Molecule - a collection of chemically bonded atoms Pure substance - one type of element, fixed chemical composition, well- defined physical and chemical characteristics physical properties - chemical properties don’t change, only the amount and physical state physical states - color, viscosity, conductivity, density, opacity, melting and boiling point States of matter: Solid - fixed shape, fixed volume, can’t be compressed, no flow, closely packed particles, vibrate a lil bit, can’t move past each other, Liquid - no fixed shape, fixed volume, can’t be compressed, has flow, loosely packed particles, don’t vibrate, can move past each other Gas- no fixed shape, no fixed volume, can be compressed, has flow, far apart particles, vibrates or moves a lot, can move past each other Physical change - change in physical properties, not chemical properties. Ex: Melting point, Freezing Point, Boiling point Intensive property - property that can’t be changed, constant. Ex: color, state of matter, melting point, boiling point, density, solubility, viscosity, ductility, conductivity, malleability Extensive property - property that can change, inconsistent Chemical properties - involves chemical change in material and substances interacting with other chemicals. chem reviewer 2  Chemical change - change of chemical composition of something. Ez: Flammability Mixture - substance made of up of two or more substances that have not chemically reacted, not a pure substance but can be separated into them, physical and chemical properties dependent on composition of mixture. Types: Homogenous - constant composition Heterogenous - no uniform composition Separating mixtures: Density - Decantation and centrifugation Boiling point - Distillation State of matter or Solubility - Filtration Intermolecular forces - Chromatography Vapor pressure - Evaporation Magnetism - Magnets Matter types: Substances: Element Compounds Mixtures: homogenous heterogenous 1.3 Units and Measurement chem reviewer 3  (see physics) Length (m) Mass (kg) - standard for mass, platinum-iridium alloy Temperature (k) - measure of how hot or cold something is Volume (L/m^3) - amount of space an object occupies Energy (J/Cal) - capacity to do work Density (g/L or g/cm^3) - relates mass to volume Significant figures - digits in a measurement, certain and uncertain, conveys precision Sig fig calcus: when add or minus, result will use least decimals when times or divide, result will use least sigfigs do pemdas precision - how close numbers are to one another accuracy - how close to exact value a number is 1.4 Unit Conversions check physics Dimensional analysis - method used to convert quantity from a unit to another Exact conversion - conversion with same system units Inexact conversion - conversion with different system of units chem reviewer 4  2.1 Development of Atomic Theory 1. Greek model - Democritus, 400 BC 2. Dalton model - 1803 3. Thomson model - 1897 4. Rutherford model - 1912 5. Bohr model - 1922 Important laws: Law of conversion of mass - matter can’t be created nor destroyed Law of constant composition - all samples of a compound has the same mass ratio Law of multiple proportions - mass ratio of elements in a compound is always a whole number Dalton’s atomic theory: (atom, same, simple, rearrange) All matter is composed of indivisible particles called matter, can’t be created nor destroyed Atoms of an element are the same, these properties are unique to each element Elements combine to form a compound with a simple ratio Atoms rearrange in chemical reactions, not created nor destroyed 2.2 Subatomic Particles Atoms have subatomic particles: Proton (p) - positive electrical charge chem reviewer 5  Electron (e-) - negative Neutron (n) - neutral Components of atom: proton and neutron found in the nucleus, small, high-density regions, electrons fill space around nucleus Neutral atom - e- = p Mass and charge of an atom - can affect physical and chemical properties of an element Atomic mass unit (u) - 1/12 the mass of a carbon atom with 6 protons and neutrons Ion - atom with unequal protons and electrons Types: Cation - atom with a positive charge, less electrons Anion - atom with negative charge, more electrons 2.3 Atoms and Isotopes Atomic number - Z = number of protons in nucleus, unique for each element, same for all atoms of the same element Neutral atom - Z = no. of e- Mass number - A = no. of p + n Atomic symbol - one or two letter symbol that represents an element, with atomic and mass no. Isotopes: atoms with same Z but different A, same p, different n named using element name and mass number chem reviewer 6  Atomic weight - avg mass of all naturally occurring isotopes of an element, takes relative abundance of isotopes into account When you consider an element’s atom mass, any sample of that element would include different isotopes with different masses Avg atomic weight = sum [(exact mass)(fractional/percent abundance)] Percent abundance - the percentage of isotope atoms in a pure element sample, describes the isotope composition of an element % abundance = no. of atoms of an isotope/total atoms of all isotopes Atomic weight(mass) - weighted avg of all element isotope Atomic mass = (%abundance of isotope 1)(isotope 1) + (%abundance of isotope 2) (isotope 2)… Isotope abundance - Mass spectrometer gives information on the mass and relative abundance of each element’s isotopes 2.4 Elements and the Periodic Table Periodic table - an important tool for chemists, contains info abt each element, organizes elements by chemical and physical properties, developed by Dmitri Mendeleev (1834-1907) Arrangement: each entry in the table has an element with Z, A, and X elements are arranged by: Groups - columns, vertical, arrange elements with repeating properties, similar physical and chem properties, has 18 groups Periods - row, horizontal, represents periodic trends or changes that occur from left to right, has 7 periods chem reviewer 7  Special group names: 1A - Alkali metals 2A - Alkaline earth metals 6A - Chalcogens 7A - Halogens 8A -Noble gases Element groups Main element groups: Group A/representative elements Transition metals: between 2A and 3A, all metals, fills B group from 4-7 period IUPAC numbers groups from 1-18 Transition elements: mostly occur in combination with other elements, except Cu, Ag, Au, and Pt cuz they’re less reactive Lanthanides and actinides - periods 6 and 7, extended portions. Lanthanides - elements between Lanthanum and lutenium Actinides - elements between actinium and lawrencium Metals - left side of the periodic table, shiny, solid, ductile, good conductor Non-metals - right side of periodic table, brittle solid or gas, bad conductors Metalloids (semimetals) - at interface of metals and nonmetals, has properties of metals and nonmetals chem reviewer 8  Group 1A - extreme left, solid at room temperature, react with air, water and halogens, found as compounds, never free elements Group 2A - metals that naturally occur as compounds, reacts with water to produce alkaline solutions, Mg(7th) Ca(5th) most abundant element Group 3A - all metals, except boron (metalloid), aluminum most abundant cuz 8.2% of earth’s mass, exceeds abundance only by nonmetal O and metalloid Si Group 4A - varies in properties cuz of change from nonmetallic to metallic behavior, element form compounds with analogous formulas, Group 5A - N occurs as diatomic molecule, 3/4th of earth’s atmosphere, P important constituent of bone, teeth and DNA, glows in the dark when exposed to air Group 6A - oxygen 20% of atmosphere, combines readily with almost all other elements, provides energy that powers life on earth Group 7A - has nonmetals, most are diatomic molecules, most reactive of all elements, combines violently with alkali metals to form salts Group 8A - least reactive elements, perceived as incapable to combine chemically to form stable compounds, until not, cuz XeF4 was prepared Elements and their states: most elements are solid, mercury liquid metal, bromine liquid nonmetal at room temperature 11 elements are gases at room temp: H, N, O, F, Cl, He, Ne, Ar, Kr, Xe, Rn How elements exist in nature: Many elements exist as molecules: diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2 Molecules consisting of two or more atoms: C60, S8, P4 Molecules connected by 3D arrangements: Si, B, Sn chem reviewer 9  Allotropes - same element that can differ in physical and chemical properties. ex: Red and white phosphorus Red P - long chain of P, red, nontoxic, burns in air above 250 degrees C White P - contains 4 P, white, yellow waxy solid, poisonous, ignites in air above 50 degrees C 2.5 The Mole and Molar Mass of Elements Avogadro’s Number = 12 g of C-12 (6.022x10^23) estimate is made using electron’s charge, can be used to convert between moles and no. of particles Molar mass of elements - mass, in g, of one mole of atoms of the element, related to atomic weight of an element, 1 mole g = amu mass units (u) 3.1 Atomic Line Spectra and the Bohr Model of Atomic Structure Niels Bohr (1885-1962): credited for the first model of the hydrogen atom, based on his explanation of the atomic line spectra. proposed electrons could occupy only certain energy levels, energy or electrons are quantized bohr’s atomic electronic structure suggests that e- moves around the nucleus in defined energy levels called orbits. electrostatic forces of e- and nucleus are balanced cuz of centripetal forces of orbiting e- Wavelength of excited gas phase elements light emitted by excited gas phase elements has unique wavelengths chem reviewer 10  e- can jump between allowed energy levels if the color(wavelength) is known, one can determine the magnitude of the energy gaps using planck’s law Bohr Model: line spectra of elements (circles around the nucleus in the Bohr model) shows that e- are placed in specific energy states called orbits principle quantum number (n) is used as a label color is produced during emission. this is when e- jumps from high to lower level orbitals Terminology: Ground state - lowest energy level of an e- Excited state - highest energy level of e- Energy levels: which energy level an e- is located is assigned by quantum number n Higher n = higher energy shell 3.2 Quantum Theory of Atomic Structure Quantum mechanics: shows a different view of how e- are arranged around a nucleus Concepts: wave behavior of matter and the uncertainty principle Wave properties of matter: chem reviewer 11  Louis de Broglie - all matter in motion has a characteristic wavelength Matter, like e-, displays wave-particle duality (properties of wave motion plus particle behavior) Heisenberg Uncertainty Principle: impossible to know with great certainty both an e-’s position and momentum(related to kinetic energy) at the same time, only probability The energy of the electron can be determined with great accuracy Schrodinger Equation and Wave functions: Erwin Schrodinger proposed that matter can be described as a wave, e- is treated as a wave and particle The wavelike behavior of e- is described by the wave function. It predicts the energy and most probable region the e- is located Boundry surface region of space where e- is likely found 3.3 Quantum Numbers, Orbitals, and Nodes Quantum numbers: resulted from Schrodinger’s wave equation for H relates the size and energy of the shell in which an orbital is placed, the shape of the orbital, and the orientation describes the location of an electron in an orbital Quantum numbers: chem reviewer 12  Each e- has their own set of quantum numbers Principle quantum number (n) - describes the energy level and size of a shell that an orbital resides in Ex: n = 1, 2, 3, 4 Angular momentum quantum number (l) - describes the orbital shape Ex: l = 0, 1, 2, 3 Magnetic quantum number (ml ) - orbital’s orientation ​ Ex: -3, -2, -1, 0, 1, 2, 3 How to find l with n given: n - 1 = l How to find ml with l given: (2l + 1) = ml  ​ ​ Orbital shapes: each boundary of l contains the volume where e- will most likely be found orbitals are classified by energy, shape, and orientation in space different regions of e- densities of an orbital are separated by nodes Electrons in Orbitals: High-energy electron: high n value, farther from the nucleus, have small negative energy Low-energy electrons: low n value, closer to the nucleus, large negative energy Types of orbitals: How to find how many orbitals are in a subshell: (2l + 1) = ml /orbitals ​ Nodes: chem reviewer 13  areas with no e- density wave energy is canceled out in nodes Types: planar and radial How to find total nodes: n - 1 How to find planar nodes: planar nodes = l How to find radial nodes: total nodes - l Energy diagrams: used to depict orbital energies Changes in Electronic State: When the ground state finds an electron in the 1s orbital, an H atom absorbs a photon of light. This results with the promotion of an electron to a higher level. The wavelength of absorbed light determines the level to which the electron is promoted 4.1 Electron Spin and Magnetism Electron and Quantum number: Electron spin - negative charged e- spinning on an axis e- only have two possible spin states: +1/2 or -1/2 the spin states are called spin quantum number when e- spin counter-clock = north is up, when e- spin clockwise = north down Symbolizing electron spin symbolized by arrows, positive is up, negative is down Magnetic materials chem reviewer 14  when an atom has an equal number of spin, the total net field is 0. materials are magnetic if e- are unpaired Categories of Magnetism: Diamagnetic - all e- are spin-paired Paramagnetic - particles with unpaired e-, and has random direction, so weak magnetism Ferromagnetic - particles with unpaired e-, and have direction, so strong magnetism 4.2 Orbital energy Orbital energy - depends on the degree of attraction between e- in that orbital and the nucleus Orbital energy of an H atom: all orbitals of a subshell are equal in energy, energy is dependent on the value of n Orbital energies of multielectron species: depends on the n and l e- energy depends on the nearness of the e- to the nucleus and the degree an e- experiences repulsion from other e- Generalizations about Orbital energies in multielectron species: when n increases, orbital energy also increases for orbitals of the same type as l increases, orbital energy increases chem reviewer 15  as n increases, subshell energies become more closely spaced and overlapping occurs 4.3 Electron Configuration of Elements Electron Configuration: shows how electrons are distributed in orbitals of an electron predicted for an atom’s ground state representation of the orbitals helps know the chemical properties and reactivity of elements Orbital filling order: Order of subshell filling is related to n and l electrons fill orbitals in order of increasing (n + l) when two or more subshell have the same (n + l) value, electrons fill the orbital with the lower n first 1s, 2s, 2p, 3s, 3p, 3s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, 8s Pauli exclusion principle - no two electrons have the same set of quantum numbers Notations: spdf notation, Ex: He = 1s2  Orbital box notation Hund’s rule - lowest-energy orbitals are half-filled first before being filled Noble gas notation, noble gas + valence electrons chem reviewer 16  Aufbau principle – Lower energy orbitals fill first Hund’s rule – Degenerate orbitals are filled with electrons until all are half-filled before electron pairing can occur Pauli exclusion principle – Individual orbitals only hold two electrons, and each should have a different spin Common traits of elements within groups: For main elements: valence e- = group number s-block = far left, p-block = far right, d-block = transitional metals, f-block = lanthanides and actinides Common traits within periods: n = period no. 4.4 Properties of Atoms Property of atoms: e- configuration plays a key role in identifying the properties of atoms factors influenced: size of largest orbital occupied energy of the highest-energy orbital number of orbital vacancies number of electrons in the highest-energy orbitals atomic properties are attributed to the degree of attraction between valence electrons and the nucleus and the no. of valence electrons High energy orbitals = lower left Low energy orbitals = upper right chem reviewer 17  Trends in Orbital energies: (more in lower left) Down within a group: energy orbital increase greater number of electrons electron orbitals with higher n larger orbitals from left to right in a period, orbital energy decreases Effective nuclear charge: (more in right) combination of attractive forces between e- and the nucleus and e- to e- repulsion How to solve: Z* = Z - (no. of core electrons) Increases from left to right across a period Atomic size: (more in the lower left) Covalent radius - the distance between nuclei of two atoms of an element when they are held together by abond Metallic radius - the distance between the nuclei of two atoms in a metallic crystal Bond distance - the distance between 2 nuclei The size of an atom is related to its electron configuration Trends in Atomic Radii: down a group, atomic radii increase and n increases, e- added are farther from the nucleus across a period from left to right, atomic radii decrease Ionization energy: (more in upper right) chem reviewer 18  amount of energy required to remove an e- from a gaseous atom values are positive increases moving left to right across a period Exceptions in IE trends: Group 3A - lower IE energy cuz np-block electrons are easier to remove than ns electrons. this is due to high energy required to remove group 3A electrons Group 6A - e- to e- (electron-electron) repulsive forces make 6A electrons easy to remove Electron affinity: (more in upper right) energy change when gaseous atoms gains an e- negative EA means energy is released large negative e- means element is more stable as an anion than as a neutral atom Trends in EA: less EA lower in a group, more upward in a group more from left to right across a period Exceptions in EA trends: Group 2A and 8A - EA close to 0, cuz in group 2A, added electrons occupy the np-block orbital, and adding electrons needs more energy Group 5A - has less EA values, cuz added electrons fills a half-filled np-block orbitals and introduces e- to e- repulsion forces 5.1 Formation and electron Configuration of Ions Coulomb’s Law: chem reviewer 19  balance of attractive and repulsive forces results in chemical bonds and decrease in energy Attractive forces have opposite charges Repulsive charges has like charges Electrostatic forces attractive forces between two oppositely charged ions: increase as ion charges increase decrease as the distance between ions increases bond distance in an ionic compound measured as a function of the sum of the individual ionic radii Cations formed when an atom loses one or more e- metals are mostly cations cuz of low IE formed by using energy, more energy is used when e- is nearer to nucleus or is equivalent to core electrons Predicting cation charge based on periodic groups: Metals in group 1A, 2A, and 3A positive charge = group number other metals are not easy to predict transition metals range from +1 to +3 besides aluminum, metals in group 3A, 4A, and 5A are not easy to predict Anion formed when an atom gains one or more e- nonmetals are mostly anions cuz of high EA formed cuz of vacancy chem reviewer 20  Lewis symbol: Amount of valence electrons around element symbol Ion size related to the size of the atom it came from and the ion formed cations are smaller cuz they lose electrons, so less space filled anions are bigger cuz they gain electrons, so more space filled Isoelectronic Ions element species with the same number of e- but different p Summary of trends: Atomic radii: more on lower left IE (ionization energy): more in upper right EA (electron affinity): more in upper right 5.2 Polyatomic Ions and Ionic Compounds Polyatomic Ions: Compounds that are ions, has names with no rules, most polyatomic ions are anions Representing Ionic Compound Symbols: Cation first before Anion, they have no net charge Nonmetal Ion Names: add “-ide” at the end of element name Polyatomic Anion Names: add “-ide”. “-ate” or “-ite” depends on how many molecules element has, more uses “-ate”, less uses “-ite” Main Metal Ion Names: add “ion” next to element name Transitional Metal Ion Names: use roman numeral to show charge then add “ion” chem reviewer 21  Ionic compound names: add previous rules, also ionic compounds balance out each, so they have a 0 net charge Polyatomic names: Starts with P: Anion: PO34 −= Phosphate ​ Starts with M: Anion: − MnO4 = Permanganate ​ Starts with O: Anions: OH− = Hydroxide OCN− = Cyanate Starts with N: Cations: Anions: NH+ 4 = Ammonium ​ NO− 2 = Nitrite ​ NO− 3 = Nitrate ​ Starts with H: Anions: HSO− 4 = Hydrogen sulfate (bisulfate) ​ HPO24 −= Hydrogen phosphate ​ H2 PO− ​ 4 = Dihydrogen Phosphate ​ chem reviewer 22  HCO− 3 = Hydrogen carbonate ​ Starts with S: Anions: SO23 −= Sulfite ​ SO24 −= Sulfate ​ S2 O23 −= Thiosulfate ​ ​ SCN− = Thiocyanate Starts with C: CN− = Cyanide CH3 CO− 2 = Acetate ​ ​ ClO− = Hypochlorite ClO− 2 = Chlorite ​ ClO− 3 = Chlorate ​ ClO− 4 = Perchlorate ​ CO23 −= Carbonate ​ C2 O24 −= Oxalate ​ ​ Cr2 O27 −= Dichromate ​ ​ CrO 2 4 −= Chromate ​ chem reviewer 23