unit 2 chem flashcards

Lesson 3.1 & 3.2 Early Atomic Theories, The Origins of Quantum Theory Vocab electron: a negatively charged subatomic particle radioactivity the spontaneous decay or disintegration of the nucleus of an atom nucleus the dense centre of an atom with a positive charge proton a positively charged subatomic particle neutron an electrically neutral subatomic particle isotopes atoms with the same number of protons but different numbers of neutrons atomic number (Z ) the number of protons in a nucleus mass number (A ) the total number of protons and neutrons in a nucleus radioisotope an isotope that emits radioactive gamma rays and/or subatomic particles (for example, alpha and/or beta particles) photoelectric effect electrons are emitted by matter that absorbs energy from shortwave electromagnetic radiation (for example, visible or UV light) quantum a unit or packet of energy (plural: quanta) photon a unit of light energy Describe the photoelectric effect. The photoelectric effect is when light shines on a metal surface and causes electrons to be ejected from it. This showed that light behaves like particles (photons) carrying energy, not just waves, and helped develop quantum theory. Explain the importance of Rutherford's gold foil experiment. In this experiment, alpha particles were directed at thin gold foil, and most passed through while some were deflected sharply. This proved that atoms have a small, dense, positively charged nucleus, replacing the earlier "plum pudding" model and forming the basis of the modern atomic model. The Development of Atomic Models What was inadequate about Rutherfors atomic model? Rutherford's atomic model could not explain the chemical properties of an element Rutherford atomic model could not explain why objects change color when heated The Bohr Model What was the new proposal in bohr model of the atom Bohr propose that an electron is found in specific circular paths or orbits around the nucleus Each possible electron orbit in Bohr’s model has a fixed energy. The fixed energies and electrons can have are called energy levels A quantum of energy is the amount of energy required to move an electron from one energy level to another energy level Like the rung of stage ladder the energy levels in an atom are not equally spaced The higher the energy level occupied by an electron, the less energy it takes to move from that energy level to the next higher energy level The quantum mechanical model What does the quantum mechanical model determine about the electrons in an atom? The quantum mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus Austrian physicist Erwin Schrodinger (1887- 1961) used new theoretical calculations and results to devise and solve a mathematical equation describing the behavior of the electron in a hydrogen atom The modern description of the electrons in atoms, the quantum mechanical model, comes from the mathematical solutions to the Schrodinger equation. In the quantum mechanical models the probability of finding an electron within a certain volume of space surrounding the nucleus can be represented as a fuzzy cloud. The cloud is more dense where the probability of finding the electron is high Atomic Orbitals How do sublevels of principal energy levels differ? An atomic orbital is often thought of as a region of space in which there is high probability of finding an electron Each energy sublevel corresponds to an orbital of a different shape which describes where the electron is likely to be found Different atomic orbitals are denoted by letters. The s orbitals are spherical, and p orbitals are dumbbell-shaped Summary of principal energy levels, sublevels and orbitals Principal Energy Level # of Sub levels Types of Sublebes n= 1 1 1s (1 orbital ) n=2 2 2s ( 1 orbital ) 2p (3 orbital) n=3 3 3s (1 orbital) 3p(3orbital) 3d (5 orbital) n=4 4 4s (1 orbital) 4p (3 orbital) 4 d (5 orbital) 4f (7 orbital) Maximum Number Energy level n Maximum number of electrons 1 2 2 8 3 18 4 32 Lesson 3.3 - Quantum Mechanical Model Vocab quantum mechanics the application of quantum theory to explain the properties of matter, particularly electrons in atoms orbital the region around the nucleus where an electron has a high probability of being found Heisenberg’s uncertainty principle the idea that it is impossible to know the exact position and speed of an electron at a given time wave function the mathematical probability of finding an electron in a certain region of space electron probability density the probability of finding an electron at a given location, derived from wave equations and used to determine the shapes of orbitals; also called electron probability distribution quantum mechanical model a model for the atom based on quantum theory and the calculation of probabilities for the location of electrons Electron Configurations What are the three rules for writing the electron configuration of element The ways in which electrons are arranged in various orbitals around the nuclei pf atoms are called electron configurations Three rules - Aufbau principle the pauli exclusion principle and hunda rules - tell you how to find the electron configuration of atoms Aufbau principle According to the aufbau principle, electrons occult the orbitals of lowest energy first. In the aufbau diagram below, each box represents an atomic orbital. Pauli Exclusion Principle According to paul expulsion principle an atomic orbital may describe at most two electrons. To occupy the same orbital two electrons must have opposite spin that is electron spins must be paired Hund’s Rule Hund’s rule states that electrons occupy orbitals of the same energy in a way that make the number of electrons with the same spin direction as large as possible Lesson 3.4 - Quantum Theory and the Electronic structure of Atoms Vocabulary Terms Quantum numbers numbers that describe the quantum mechanical properties of orbitals; from the solutions to Schrödinger’s wave equation Shell an atom’s main energy level, where the shell number is given by the principal quantum number, n = 1, 2, 3, … Principal quantum number (n ) the quantum number that describes the size and energy of an atomic orbital Subshells orbitals of different shapes and energies, as given by the secondary quantum number; often referred to as s, p, d, and f Secondary quantum number (l ) the quantum number that describes the shape and energy of an atomic orbital, with whole-number values from 0 to n – 1 for each value of n Magnetic quantum number (ml ) the quantum number that describes the orientation of an atomic orbital in space relative to the other orbitals in the atom, with whole-number values between +/ and -/, including 0 Spin quantum number (ms) the quantum number that relates to the spin of the electron; limited to +1/2 or -1/2 Pauli exclusion principle the principle Quantum numbers are used to differentiate between electrons In quantum theory, each electron in an atom is assigned a set of four quantum numbers Three of these gives the location of the electron, and the fourth gives the orientation of the electron within the orbital Definitions of numbers 4 Quantum, numbers N → L → n -1 M1 Ms → + ½ , -½ Ms stands for the spin of the electrons (spins right or spins left) Quantum numbers (n, l, ml, ms) Quantum number n N = 1, 2 , 3 , 4 … ( never be zero) Distance of electrons from nucleus The further/higher you get, the more rings there will be (n=1, n=2, n=3 etc) Quantum number l : Angular momentum quantum number l For a given value of n, l = 0, 1, 2 , 3… n - 1 n=1 l = 0 L = n-1 L = 0→ s orbital L = 1→ p robistal L = 2 → d orbital L = 3 → f orbital Shape of the “volume” of space that the electron occupies L values S = 0 P = 1 D = 2 F =3 G =4 Ml values 1 room 3 rooms -1, 0, 1 5 rooms -2, -1, 0, 1, 2 7 rooms -3, -3, -1 , 0 , 1 , 2 , 3 9 rooms -4, -3, -2, -1, 0 , 1, 2, 3, 4 Quantum number ml: Magnetic quantum number ml For a given value of l Ml = -1…. 0, …. +1 If l = 1 (p orbital), ml = -1, 0 or 1 If l = 2, (d orbital) ml = -2, -2, 0, 1, or 2 “Orientation of the orbital in space” Ml = -1, 0 or 1 → 3 different orientations is space Ml = -2, -1, 0, 1 or 2 → 5 orientations is space Quantum number ms Spin quantum number ms Existence (and energy) of electron in atom is described by its unique wave function Ψ Pauli exclusion principle - no two electrons in an atom can have the same four quantum numbers Each orbital can hold a certain number of levels Quantum numbers (n, l, ml, ms) Shell - electrons with the same value of n Subshell - electrons with the same values of n and l Orbital - electrons with the same values of n, l and m How many electron can and orbital hold Each orbital takes a maximum of 2 electrons Is n, l and ml are fixed, then ms = ½ or -½ Ψ = (n, l , ml ½) or Ψ = (n, l, ms, -½) N = 2 → 2p → l = 1 If l = 1, then ml= -1, 0, or +1 3 orbitals The energies of orbitals Energy of orbitals in a single electron atom (orbital diagram) Energy only depends principle quantum number n Energy of orbitals in a multi electron atom Energy depends on n and l “Fill up” electrons in lowest energy orbitals (Aufbau principle) First electron has to be very close to the nucleus, when that energy level is filled up, then you go to the next one. The most stable arrangement of electron in subshells is the one with the greatest number of parallel spins (Hund’s rule) VERY IMPORTANT TO KNOW HOW TO DRAW ORBITAL DIAGRAMS Electron configuration is how the electrons are distributed among the various atomic orbitals in an atom 1s^1 1 = principle quantum number m S = angular momentum number 1 ^1 number of electrons in the orbital of subshell What is the electron configuration of Mg? Mg 12 electrons 1s < 2s < 2p < 3s < 3p < 4s 1s^2 2s^2 2p^6 3s^2 2 + 2 + 6 + 2 = 12 Abbreviated as [Ne] 3s^2\ What are the possible quantum number for the last (outermost ) electron in Cl?Lesson 3.5 - Atomic structure & Periodic table electron configuration: the location and number of electrons in the electron energy levels of an atom aufbau principle: the theory that an atom is “built up” by the addition of electrons, which fill orbitals starting at the lowest available energy orbital before filling higher energy orbitals (for example, 1s before 2s） energy-level diagram (orbital diagram): a diagram that represents the relative energies of the electrons in an atom Hund’s rule: a rule stating that in a particular set of orbitals of the same energy, the lowest energy configuration for an atom is the one with the maximum number of unpaired electrons allowed by the Pauli exclusion principle; unpaired electrons are represented as having parallel spins valence electron: an electron in the outermost principal quantum level of an atom representative elements: those elements in the main blocks of the periodic table, which are Groups 1 to 18 (the s and p blocks) Ferromagnetism: the very strong magnetism commonly exhibited by materials that contain nickel, iron, and cobalt; applies to a collection of atoms paramagnetism：the weak attraction of a substance to a magnet; applies to individual atoms Electron configurations What are the three rules for writing the electron configurations of elements? The ways in which electrons are arranged in various orbitals around the nuclei of atoms are called electron configurations. Three rules - the aufban principle, the pauli exclusion principle, and Hund’s rule- tell you how to find the electron configurations of atoms Exceptional electron configurations Why do actual electron configurations form some element differ from those assigned using the aufbau principle Lesson 4.1 - Chemical bonds Ionic bond - the electrostatic attraction between oppositely charged ions Isoelectronic - having the same number of electrons per atom, ion, or molecule Covalent bond - a chemical bond in which atoms share the bonding electrons Bonding electron pair - an electron pair that is involved in bonding, found in the space between 2 atoms Lewis structure - a diagram that represents the arrangement of covalent electrons and bonds in a molecule or polyatomic ion Duet rule - the observation that the complete outer shell of valence electrons when hydrogen and period 2 metals are involved in bonding Octet rule - the observation that many atoms tend to form the most stable substances when they are surrounded by 8 electrons in their valence shells Lone electron pair a pair of valence electrons that is localized to a given atom but not involved in bonding Simplified Lewis structure - a Lewis structure in which bonding electron pairs are represented by solid lines and lone electron pairs by dots Space-filling model - a model of a molecule showing the relative sizes of the atoms and their relative orientations Coordinate covalent bond/Dative bond - a covalent bond in which the electrons involved in bonding are from one atom Chemical bond: Atoms or ions strongly attached to one another There are three types: Ionic bonds Covalent bonds Metallic bonds Ionic bonds Electrostatic force that exists between particles of opposite charge that results from a transfer of electrons metals to non metals Common Features of Ionic Bonds Ionic bonds form metals and non-metals In naming simple ionic compounds the metal is always first, the non metal second (i.e sodium chloride) In solution, ionic compounds easily conduct electricity Ionic compounds ionize easily in water and other polar solvents Ionic compounds tend to form crystalline solids with high melting points Covalent's bond Sharing of electrons between two non metals Sharing can be equal (non polar) Sharing can be not equal (polar) Features of covalent bond Each atom shares its unpaired electron, both atoms are “tricked” into thinking each has a full valence of eight electrons Tend to be gases, liquids or low melting point solids, because the intermolecular forces of attraction are comparatively weak Most covalent substances are insoluble in water but are soluble in organic solutions Most not all, example is sugar can dissolve in water because it is polar Poor conductors Metallic Bonds Bonds between metals Metallic bonds don't actually bond with each other, its just the attraction between the ions Metal have low ionization energies thus they do not have a tight hold on their valence electrons Thus formicin and electron sea that cements the positives nuclei together and shields the positive cores from each other The electrons are not bound to anyone particular atom and are free to move when an electrical field is applied This accounts for the electrical conductivity of metals, and also their thermal conductivity since the moving electrons carry thermal vibration energy from place to place as they move. Features of metallic bonds Metals are good conductors of heat and electricity. This is directly due to the mobility of the electrons The “cement” effect of the electrons determines the hardness of the metal. Some metals are harder than others, the strength of the “cement” various from metal to metal Metals are lustrous (shine) Metals are malleable (can be flattened) and ductile (can be drawn into wires) because of the way the metal cations and electrons can “flow” around each other without breaking the crystal structure. Valence Electrons The electrons in the outer most shell of an atom that are involved in bonding The number of valence electrons an atom has is the group number E.g group 1A or IA = 1 valence electron Lewis structures A method used to illustrate valence electrons and bonding between atoms. E.g Sulfur = group 6 = 6 valence e Octet rules Rules of eighth Atoms tend to gain share or lose electrons until they have 8 electrons in their valence shell Note what the largest group number is Exception: Hydrogen = Rule of 2 Ionic bonding Look at the balanced reaction of sodium (metal) and chloride (non metal) Na (s) +1/2Cl2(g) → NaCl (s) Note: ΔHf = -410.9 kJ Therefore we have an enthalpy change that is exothermic (exo = out) Illustrating covalent bones each pair of shared electrons is a line C-C Unshared electrons are dots Multiple bonds Single bond 2 atoms share 1 pair or electrons C-C Double bond 2 atoms share 2 pairs of electrons C=C Triple bond 2 atoms share 3 pairs of electrons C≡C Bond lengths and strength In general as the number of bonds between two atoms increases the bond length grows shorter and stronger. Lesson 4.2 - Bonding Theories Vocab Three dimensional structures : the 3D arrangements of ions or atp,s making ip a pir substance Valence she electron pair repulsion (VESPR) theories : a method to determine the geometry of a molecule based on the idea that an electron pairs are as dark paper as possible Electron pair repulsion: the repulsive force that occurs between electron pairs causing them to be positioned as far apart as possible in molecules VSEPR theory How does VSEPR theory help predict the shapes of molecules? The valence shell electron pair resolution theory or VSEPR theory, explains the three dimensional shape of methane According to VSEPR theory, the repulsion between electron pairs causes molecular shapes to adjust so that the valence electron pairs stay as far apart as possible. The measured bond angle in water is about 105 degrees Nine possible molecular shapes Linear Triatonic Tionbal planer Bent triatomic Puramida Tetrahedral Trigonal Bipyramidal Octahedral Square planar T-shaped Hybrid Orbitals In what ways is orbital hybridization useful in detecting molecules? Orbital hybridization provides information about both molecules bonding and molecular shape In hybridization several atomic orbitals mix to form the same total number of equivalent hyderabad orbitals. SP(1)- Hurpidazation Hybridization involving single bonds Atomic orbitals of two hydrogen atoms Hybrid orbitals of a carbon atom Atomic orbitals of two hydrogen atoms Hybridization involving double bonds Hybridization involving triple bonds Shapes of molecules Molecules (covalent chemicals) form certain shapes depending on how many lone and bonding pairs of electrons it has Because the electron pairs repel each other we get certain shapes being formed. These are due to a certain rule called VSEPR. VSEPR theory Based on lewis structures we can know the shape or “geometry” of molecules VSEPR, as the name suggests, predicts geometry based on the repulsion of electron pairs (Bonding pairs and lone pairs). Electrons around the central nucleus repel each other. Thus, resulting structures have atoms maximally spread out. AXE Lewis structures do not show geometry, only electron pair placement. However, the 3d shape (geometry) of a molecule can be determined from a properly drawn lewis structure All monocentric molecules can be represented by an AXE formula A = central atom X = outer atoms (doesn't matter what they actually are or how many bonds they are held by) E = lone pairs of electrons on the central atom only What AXE formula correspond to sulfur trioxide, SO3 Lesson 4.3 - Electronegativity and bond polarity Vocabulary non-polar covalent bond a covalent bond in which the electrons are shared equally between atoms polar covalent bond a covalent bond in which the electrons are not shared equally because 1 atom attracts them more strongly than the other atom electronegativity the ability of an atom in a molecule to attract shared electrons to itself dipole a separation of positive and negative charges in a region in space dipole dipole an intermolecular force that occurs between polar molecules, where the partially positive end of one molecule is attracted to the partially negative end of another Electron affinity the energy change that occurs when an electron is added to a neutral gaseous atom to form a negative ion Bond polarity How do electronegativity values determine the charge distribution in a polar bond? When the atoms in a bond pull equally (as occurs when identical atoms are bonded), the bonding electrons are shared equally, and the bond is a nonpolar covalent bond The bonding pairs of electrons in covalent bonds are pulled by the nuclei The chlorine atom attracts the electron cloud more than the hydrogen atom does A polar covalent bond known as a polar bond, is a covalent bond between atoms in which the electrons are shared unequally The more electronegative atom attracts electrons more strongly and gains a slightly negative charge. The less electronegative atom has a slightly positive charge. Polar molecules: What happens to polar molecules between a pair of oppositely charged metal plates? In a polar molecule, one end of the molecule is slightly negative, and the other end is slightly positive. A molecule that has two poles is called a dipolar molecule, or a dipole When polar molecules are placed between oppositely charged plates, they tend to become oriented with respect to the positive and negative plates A hydrogen chloride molecule is a dipole Attractions between molecules How do intermolecular attractions compare with ionic and covalent bonds? Intermolecular attractions are weaker than ionic or covalent bonds. These attractions are responsible for determining whether a molecular compound is a gas, liquid, or a solid at a given temperature. Van Der Waals forces The two weakest attractions between molecules are collectively called van der waals forces. Named after the Dutch chemist Johannes van der Waals (1837-1923) Dipole interactions occur when polar molecules are attracted to one another Dispersion forces, the weakest of all molecular interactions, are caused by the motion of electrons. The strength of dispersion forces generally increase as the number of electrons in a molecule increases Hydrogen bonds Hydrogen bonds are attractive forces in which a hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another electronegative atom. The relatively strong attractive forces between water molecules cause the water to form small drops on a waxy surface. Intermolecular attractions and molecular properties Why do network solids have high melting points? Network solids (or network crystals) are solids in which all of the atoms are covalently bonded to each other Network solids consist of molecules that do not melt until the temperature reaches 1000 degrees celsius or higher, or they decompose without melting at all. Diamond is an example of a network solid. Diamond does not melt. It vaporizes into a gas at 3500 degrees celsius or above. Silicon carbide is a network solid. It has a melting point of about 2700 degrees celcius. Types of bond Ionic > 2.0 Polar covalent 0.5 - 2.0 Nonpolar covalent < or equal to 0.5 For each pair, Calculate bond polarities Identify each bond as ionic, polar covalent, or non polar covalent. HBr ΔEN = 2.1 – 2.8 = 0.7 → polar covalent H₂O ΔEN = 3.5 – 2.1 = 1.4 → polar covalent LiI ΔEN = 1.0 – 2.8 = 1.8 → ionic BrCl ΔEN = 2.8 – 3.0 = 0.2 → nonpolar covalent NH₃ (N–H bond) ΔEN = 3.0 – 2.1 = 0.9 → polar covalent KF ΔEN = 1.0 – 4.| = 3.0 → ionic Lesson 4.8 The Structure & Properties of Solids Vocb composite material a material composed of two or more distinct materials that remain separate from each other in the solid phase metallic crystal a solid with closely packed atoms held together by electrostatic interactions and free-moving electrons electron sea theory a theory that states that the electrons in a metallic crystal move freely around the positively charged nuclei metallic bonding the bonding that holds the nuclei and electrons of metals together molecular crystal a solid composed of individual molecules held together by intermolecular forces of attraction covalent network crystal a solid in which the atoms form covalent bonds in an interwoven network carbon nanotube a solid made of carbon atoms similar to graphite rolled into a cylinder buckyball a spherical arrangement of carbon atoms that forms a hollow, cage-like structure semiconductor a substance that conducts a slight electric current at room temperature but has increasing conductivity at higher temperatures A model for solids How are the structure and properties of solid related? The general properties of solids reflect the orderly arrangement of their particles and the fixed locations of their particles. The four types of solids Ionic Molecular Covalent Metallic The melting point is the temperature at which a solid changes into a liquid Crystal structure and unit cells What determines the shape of a crystal? In a crystal, the particles are arranged in an orderly, repeating, three dimensional pattern called a crystal lattice The shape of a crystal reflects the arrangement of the particles within the solid Crystal shapes A crystal has sides, or faces. Crystals are classified into seven crystal systems Seven structures: Hexagonal Triclinic Rhombohedral Monoclinic …….. The smallest group of particles within a crystal that retains the geometric shape of the crystal is known as a unit cell A crystal lattice is a repeating array to any one of 14 kinds of unit cells There are from one to four types of unit cells that can be associated with each crystal system Three kind of units cells can make up a cubic crystal system Simple cubic The atoms or ions are arranged at the corners of the imaginary cube Body centered The atoms or ions are at the corners and in the center of the imaginary cube Face centered The atoms or ions at the corners and in the center of each face of the imaginary cube Allotropes Allotropes are two or more different molecular forms of the same element in the same physical state Carbon Allthropes Diamond In diamond each carbon atom in the interior of the diamond is strongly bonded to four others the array is rigid and compact Graphite In graphite the carbon atoms are linked in widely spaced layers of hexagonal (six-sided) arrays Fullerene In buckminsterfullerene, 60 carbon atoms form a hollow sphere. The carbons are arranged in pentagons. and hexagons Non crystalline solids An amorphous solid lacks an ordered internal structure Rubber, plastic, asphalt, and glass are amorphous solids A glass is transparent fusion product of inorganic substance that have cooked to a riginf state without crystallizing