G11 Chemistry Unit Test 2 on Atomic Structure and Periodicity Reviewer PDF

Reviewer: Atomic Structure: Introduction to Atoms: Atom: The basic unit of an element that can enter into a chemical combination. In real life, the shape of an atom is not a simple geometric figure but is described by probabilistic electron cloud models. Basic Laws of Matter: 1. Law of Conservat...

Reviewer: Atomic Structure: Introduction to Atoms: Atom: The basic unit of an element that can enter into a chemical combination. In real life, the shape of an atom is not a simple geometric figure but is described by probabilistic electron cloud models. Basic Laws of Matter: 1. Law of Conservation of Mass: In a closed system, the total mass of substances involved in a chemical reaction remains constant, regardless of the changes that occur during the reaction. 2. Law of Constant Composition: A given chemical compound always contains its component elements in fixed ratio by mass and does not depend on its source or how it was prepared. 3. Law of Multiple Proportions: When two elements form more than one compound between them, the different masses of one element that combine with a fixed mass of the other element are in simple, whole-number ratios. Atomic Theory: 1. All matter is composed of tiny indestructible particles called atoms. 2. Atoms cannot be created or destroyed. 3. Atoms of the same element are alike in every way. 4. Atoms of different elements are different. 5. Atoms can combine together in small numbers to form molecules. Nuclear Atom Symbol Notation: A ZX A: Mass Number (# of protons + # of neutrons) Z: Atomic Number (# of protons) X: Chemical Symbol of an element Counting Sub-atomic Particles: For Neutral Atoms: # of protons = atomic number (Z) # of protons (p) = # of electrons (e) # of neutrons (n) = mass number (A) – atomic number (Z) mass number = # of protons + # of neutrons Ions and Isotopes: Ion: An ion is a charge-carrying atom. It contains a different number of protons or electrons. It is the result of transferring of electrons in a chemical reaction. Cation: Positively charged. It has more protons than electrons. It loses electrons from positive ions. Anion: Negatively charged. It has more electrons than protons. It gains electrons from negative ions. Counting Sub-atomic Particles: For Ions: # of protons = atomic number (Z) # of protons (P) ≠ # of electrons (e) protons + electrons = overall charge of the ion Isotopes: Atoms of the same element that have different numbers of neutrons. It can be classified as: Stable and Unstable. Stable Isotopes: The atom’s nucleus is stable. It has stable nuclei and it does not show radioactivity. Unstable Isotopes: The atom’s nucleus is unstable. It has unstable nuclei and it shows radioactivity. Radioactivity: The release of energy from the decay of the nuclei of certain kinds of atoms and isotopes. Isotope Symbol: A ZX A: Mass Number (# of protons + # of neutrons) Z: Atomic Number (# of protons) X: Chemical Symbol of an element Isotope Percentage Abundance: The percentage of atoms with a specific atomic mass found in a naturally occurring sample of an element is known as its relative abundance. Formula: (M1)(x) + (M2)(1-x) = M(E) Where M1 – mass of first isotope M2 – mass of second isotope x – relative abundance M(E) –atomic mass of the element Relative Average Atomic Mass: Mass of an element that is affected by the individual mass of each isotope and its respective abundances. Its unit is amu (atomic mass unit). Formula: A = ∑(mass of isotope x % abundance) R 100 Mass Spectrometry: Mass Spectrometer: It measures the masses of the different isotopes and their relative abundance. Mass Spectra: It is the result of the analysis by the mass spectrometer. Vertical axis: It shows the percentage abundance. Horizontal axis: It shows the mass number. Electromagnetic Radiation: Electrons orbit in the nucleus. Energy of the orbit is related to its size. Electrons can move between each shell. (Picture on the left shows a Ground State, on the right shows an Excited State) An atom’s electron absorbs energy and becomes energized, or excited. The excited electron moves from its ground state to an excited state. The excited state produced is unstable. Electrons soon fall back to the ground state. The excited electron emits the energy in the form of electromagnetic radiation or Light. Energy is in a “package” form known as Photons. Electromagnetic Radiation: The emission and transmission of energy in the form of electromagnetic waves. The EM waves are characterized by wavelengths and wave frequencies. Wavelength: The distance between two corresponding points on adjacent waves. Frequency: The number of waves that pass a fixed point in a given amount of time. Wave Speed ( c ): The distance travelled by the wave profile per unit time. Its unit is expressed in ms -1. Electron magnetic waves travel at a speed of 3x10 m/s. Electromagnetic Spectrum: The electromagnetic spectrum encompasses all types of electromagnetic radiation. Each wavelength in the visible range produces a perception of color in the eye. Continuous Spectrum: All wavelengths of visible light are represented in the spectra. Line Spectrum: Light emission only at specific wavelengths. Line Emission Spectra of Excited Atoms: Excited atoms emit light of only certain wavelengths. The wavelengths of emitted light depend on the element. Each element produces a unique set of spectral lines. Since no two elements emit the same spectral lines, elements can be identified by their line spectrum. Quantum Numbers: Heisenberg Uncertainty Principle: It states that it is impossible to know simultaneously both the momentum and the position of a particle with certainty. The Schrödinger Equation: It describes the energy and position of electrons in an atom. Quantum Numbers: Describes the distribution of electrons in the atomic orbitals. Each electron in an atom is described by four different quantum numbers. Atomic Orbitals: The wave function of an electron in an atom. Each orbital can hold up to two electrons. 1. Principal Quantum Number (n): (n = 1, 2, 3, etc.) n = the energy level (shell) of the electron 2. Angular Momentum Quantum Number (ℓ): (ℓ = 0,..., n-1) ℓ = the shape of an orbital (s, p, d, f) 3. Magnetic Quantum Number (mℓ): (m = - ℓ,..., 0,..., + ℓ) m = the individual orbital which holds the electrons. 4. Spin Quantum Number (ms): (s = +½ or -½ ) s = the spin axis of an electron An electron can spin in only one of two directions (sometimes called up and down). Every electron in an atom has unique set of quantum numbers. No two electrons in an atom can have the same four quantum numbers as stated by the Pauli exclusion principle. Energy Level Diagram: A graphical representation of the energies of the various orbitals of an atom. Box = atomic orbital Group of boxes = subshell Arrow = electron Electrons all want to go to the lowest level. At most two electrons can occupy one atomic orbital. Two electrons must have opposite electron spin (spin up and spin down). Filling Rules for Electron Orbitals: Pauli Exclusion Principle: It states that no two electrons in an atom can share the same four quantum numbers. Hund’s Rule: It states that all orbitals in a sublevel must be half-filled with electrons having +½ spin before the orbitals can be completely filled. Aufbau Principle: It states that electrons fill orbitals of lowest energy first. Electron Configuration: Electron Configuration: It describes where electrons are located around the nucleus of an atom. 1s1 1 - Energy Level s - # of electrons in the orbital 1 - Type of orbital s – sharp p – proximal d – diffuse f – fundamental Condensed Form: A way of abbreviating long electron configurations. Electron Configuration for Ions: Cations are formed by removing electrons from atoms. Anions are formed by adding electrons to atoms. Ions of s and p blocks elements are isoelectronic (having the same of number of neutrons) with a noble gas but contain a different number of protons and are charged. Periodicity: Periodic Table, Effective Nuclear Charge, Electron Shielding: Periodicity: Regular periodic variation properties of elements with atomic number and position in the periodic table. Cause of Periodicity: Periodic repetition of similar electronic configuration of elements as the atomic number increases. Periodic Table: A chart in which elements having similar chemical and physical properties are grouped together. Elements are arranged in increasing atomic numbers. Metal - It is a good conductor of heat and electricity. Nonmetal - It is a poor conductor of heat and electricity. Metalloid - It has properties that are intermediate between those of metals and nonmetals. Periodic Trends Terms: Effective Nuclear Charge: It is the measurement of attractive force between protons and electrons. Electron Shielding: The core electrons shield the valence electrons from the full attractive forces of the protons in the nucleus. Periodic Trends (Physical Properties): Periodic Trends Terms: Effective Nuclear Charge: It is the measurement of attractive force between protons and electrons. Electron Shielding: The core electrons shield the valence electrons from the full attractive forces of the protons in the nucleus. Atomic Radius: The distance from the nucleus of an atom to the outermost electron. Atomic Radii Trend: The size of an atom or ion is determined by valence electrons. The larger the effective nuclear charge for these valence electrons, the more strongly they are drawn towards the nucleus. Decreases across the period: There is an increase in effective nuclear charge and weaker shielding effect. Increases going down: There is an increase in principal energy levels and higher shielding effect. Ionic Radius: The radius of a cation or an anion. Removing one or more electrons from an atom reduces electron-electron repulsion but the nuclear charge remains the same, so the electron cloud shrinks. If the atom forms an anion, its size (radius) increases, because the the nuclear charge remains the same but the repulsion enlarges the domain of the electron cloud. Ionization Energy: The amount of energy required to remove an electron from a gaseous atom. First Ionization Energy: Atoms in the gas phase are uninfluenced by its neighbors and so there are no forces between molecules. Ionization Energy Trend: Increases across the period: There is an increase in effective nuclear charge and constant shielding effect, more energy needed. Decreases going down: There is an increase in principal energy levels and higher shielding effect, less energy needed. Electron Affinity: The amount of energy released when an atom in the gas phase gains an electron. Electron Affinity Trend: Increases across the period: There is an increase in effective nuclear charge, decreasing atomic radius and potential stable electron configuration. Decreases going down: There is an increase in principal energy levels and higher shielding effect. Electronegativity: The ability of an atom to attract towards itself the electrons in a chemical bond. Increases across the period: There is an increase in effective nuclear charge and stronger attraction for electrons in bonds. Decreases going down: There is an increased distance from the nucleus and electron shielding. Very best of luck with the exam :)