Atoms and Atomic Structure PDF

UNIT 1 1.2 ELEMENT: a substance that cannot be broken down into simpler substances by a chemical reaction -​ all matter is built from these elements ATOM: the smallest possible particle/unit of an element, Greek “not able to be cut” Democritus and Aristotle (2500 years ago) -​ philosophers o...

UNIT 1 1.2 ELEMENT: a substance that cannot be broken down into simpler substances by a chemical reaction -​ all matter is built from these elements ATOM: the smallest possible particle/unit of an element, Greek “not able to be cut” Democritus and Aristotle (2500 years ago) -​ philosophers of ancient Mediterranean and northern Africa Democritus – atomos means “uncuttable”; atoms have different sizes, are in constant motion, and are separated by empty space Aristotle – Four Element Model accepted for 2000 years; fire, air, earth, water the Alchemists (1st c. – 17th c.) -​ explored the nature of matter intensely -​ searched for elusive prizes like the “elixir of life” (hope of immortality), the philosopher’s stone (thought to transform common metals to gold) -​ constructed lab glassware and equipment -​ developed alloy, handling procedures for dangerous chemicals and various chemical processes -​ discovered new elements such as: arsenic (As), antimony (Sb), bismuth (Bi), phosphorus (P) Dalton’s Billiard Ball Model (1808) -​ experiment: broke down water using an electric current -​ observations: hydrogen gas and oxygen gas formed; had different properties from water -​ conclusion: water is not an element but a compound made of elements 1. All matter is made of tiny indivisible atoms (“small, hard, indestructible sphere”), which are particles too small to see. 2. Atoms of the same element are identical in mass and size. Atoms of one element are different in mass and size from the atoms of other elements. 3. Compounds are created when atoms of different elements link together in definite proportions. 4. Chemical reactions involve the rearrangement of atoms. Atoms cannot be created, destroyed, or subdivided in chemical changes. Thomson’s Plum Pudding Model (1904) -​ experiment: applied high voltage to a partially evacuated tube with a metal electrode at each end -​ observations: (1) a ray was produced that started from the negative electrode (cathode) – called tube “cathode ray tube” (2) negative pole of an applied electric field repelled the ray -​ conclusion: ray was composed of a stream of negatively charged particles (electrons) 1.​ Atoms contain particles called electrons (negatively charged subatomic particles). 2.​ Electrons have a small mass (m) and a negative charge (e). 3.​ The atom is a sphere of positive charge. 4.​ Electrons are embedded in this sphere, so that positive and negative charges balance and overall the atoms is neutral or uncharged. Rutherford’s Nuclear Model (1911) -​ experiment:positively charged alpha particles were fired at a very thin sheet of gold foil -​ source of alpha particles was the radioactive element radium, fluorescent screen lit up when struck by alpha particle -​ prediction: most (+) particles pass through foil in a straight line through spaces between particles in the foil -​ results: most particles did pass through but some were repelled and deflected backwards at various angles -​ conclusion: 1. atom is mostly empty space 2. atom has a small, dense, positive (+) centre called the nucleus - (that repels and deflects alpha particles if they get too close) 3. electrons are separate from the nucleus - electrons revolve around the nucleus at a relatively far distance (like a planet around the Sun) Chadwick – Discovered neutron(1932) -​ experiment: worked with Rutherford to determine masses of nuclei of different elements -​ observations: observed masses of nuclei were not the same as the sum of the masses of the protons -​ conclusion: nucleus must contain not only positively charged protons, but also neutral (uncharged) particles called neutrons Bohr’s Planetary Model (1911) experiment: worked with applying electricity and thermal energy to hydrogen gas; directed light through a prism with a screen behind it observations: hydrogen atoms emitted light when they were “excited” by the additional energy; observed lines of only certain colours of light (atomic spectra) conclusion:electrons orbit the nucleus of the atom in definite energy levels/orbits - each energy level can hold a certain number of electrons (2, 8, 18, etc.) - when electrons jump from a higher energy level to a lower energy level they emit a certain quantity of energy BOHR-RUTHERFORD DIAGRAMS show the number of each type of subatomic particle in a specific atom and to represent the arrangement of electrons around the nucleus especially useful for the first 20 elements concentric circles represent the different energy levels the electrons are drawn in the appropriate locations, and the numbers of protons and neutrons are indicated in the nucleus Bohr model is a simplified one but it is very useful for making connections between atomic structure of different elements and their chemical and physical properties VALENCE SHELL/VALENCE ELECTRON -​ the outermost energy level or orbit of an atom or ion, an electron in the outermost energy level or orbit Bohr’s work showed that electrons release burst of light energy. The wavelength (color) of the light depends on where the electrons are jumping from and to. Hydrogen atoms, shown here, can emit four colors of light. Each element emits a unique spectrum. Schrodinger’s Quantum Mechanical Model (1926) -​ does not have well- defined orbits for electrons -​ describes electron position in terms of a ‘cloud of probability’ rather than a specific orbital path -​ does not treat the electron as a localized particle but gives a probability wave description -​ the electron has the greatest probability of being found where the cloud is darkest and the least probability of being found where it is lightest -​ predicts the relative intensities of various spectral lines 1.3 STABLE OCTET -​ an electron arrangement where the valence shell is filled with 8 valence electrons (2 for hydrogen and helium) -​ true for first 18 elements -​ noble gases (group 18) have 8 electrons in outer shell OCTET RULE a generalization stating that when atoms combine, they tend to achieve 8 valence electrons three possible ways in which an atom can achieve stable arrangement: share, lose, or gain electrons ION a charged entity formed when an atom gains or loses one or more electrons whether an atom gains or loses electrons depends on the number of valence electrons it has metals with 1, 2, or 3 valence electrons tend to lose them to form ions with charges of +1, +2, +3 respectively nonmetals with 5, 6, or 7 valence electrons tend to lose them to form ions with charges of –3, –2, –1 respectively when drawing a Bohr-Rutherford diagram for an ion, place square brackets around it and indicate the charge CATION e.g. Na+ ions are called sodium ions -​ a positively charged ion formed by the removal of one or more electrons from the valence shell of a neutral atom -​ metal atoms have few valence electrons, so they lose electrons to form cations with the electron arrangement of the nearest noble gas with a smaller atomic number ANION e.g. the Cl– ion is called a chloride ion a negatively charged ion formed by the addition of one or more electrons to a neutral atom non-metal atoms have almost complete valence shells so they gain electrons to form anions with the electron arrangement of the nearest noble gas with a larger atomic number VALENCE the charge of an ion the combining capacity of an atom determined by the number of electrons that it will lose, add, or share when it reacts with other atoms MULTIVALENT the property of having more than one possible valence most of the transition metals are multivalent classical naming system was used for naming multivalent elements with two different possible valences: used the Latin name of the element and a suix of either –ous to represent the lower valence or –ic for the higher valence IUPAC system uses a Roman numeral in the ion’s name to indicate the charge of the ion: copper(I), copper(II) POLYATOMIC ION an ion, made up of more than one atom, that acts as a single entity i.e. a polyatomic ion behaves just like an ion made of only one atom 1.4 SUBATOMIC PARTICLES  -​ nucleus is very small compared to size of atom -​ protons and neutrons are found in the nucleus and are called nucleons -​ protons are responsible for the element’s identity -​ neutrons have virtually same mass as proton electrons constitute nearly all of the volume of an atom, but an insignificant amount of its mass atoms have different chemical properties due to number of electrons and arrangement of electrons electrons interact with electrons of other atoms when they combine to form compounds ISOTOPE a form of an element in which the atoms have the same number of protons (same atomic number, Z) as all other forms of that element, but a different number of neutrons (different mass number, A) scientists use mass number to distinguish between different isotopes for a given element share chemical properties (behave the same way during chemical reactions) but their physical properties (e.g. melting point, boiling point - heavier isotopes move more slowly at a given temperature) vary different elements have different numbers of isotopes STANDARD ATOMIC NOTATION Atomic Number (Z): -​ the number of protons in the nucleus of an atom -​ determines the identity of the element -​ also represents the number of electrons in an electrically neutral atom Mass Number (A): -​ the sum of the number of protons and neutrons (nuclear particles; nucleons) in the nucleus of an atom -​ mass of atom depends on protons and neutrons only Neutron Number (N): – the number of neutrons in the nucleus of an atom –N=A–Z Ions and Standard Atomic Notation for ions include the charge on the upper right side of the symbol RELATIVE/AVERAGE ATOMIC MASS the weighted average (taking into account relative abundance) of the masses of all the naturally occurring isotopes of an element compared to a scale on which one carbon-12 atom has a mass of exactly 12 units ATOMIC MASS UNIT, u 1/12 the mass of a carbon-12 atom -​ used because atoms are extremely small) and it would be very cumbersome to represent the mass of atoms using the kg -​ atomic masses can now be measured directly by using a mass spectrometer but before the mass spectrometer this was impossible because the mass of a single atom is very tiny masses of all other atoms are measured relative to the mass of carbon-12 C-12 isotope of carbon was chosen as standard in 1961 to compare all other masses to because – C-12 is most abundant isotope of carbon – very common – easy to transport and store (solid) ISOTOPIC ABUNDANCE -​ the percentage of an isotope in a sample of an element -​ most elements exist naturally as a mixture of several isotopes -​ for most elements the isotopic composition of any sample is stable/consistent -​ mass spectrometers can determine the atomic mass and relative abundance of each isotope MASS SPECTROMETER a measuring instrument used to determine the mass and abundance of isotopes can provide information about structure of molecules composed of 3 main sections: the ion source, the analyzer, and the detector Radioisotope -​ isotope that has an unstable nucleus which spontaneously decays emitting radioactive gamma rays and/or subatomic particles (alpha and/or beta particles) -​ produces two or more smaller nuclei and radiation -​ applications in carbon dating, nuclear energy, medicine Radioactive -​ many isotopes with atomic numbers higher than 20 have unstable nuclei -​ they are said to be RADIOACTIVE - they emit radiation in order to obtain a more stable state -​ elements with atomic numbers of 84 or greater consist only of radioactive nuclei NUCLEAR RADIATION -​ energy (gamma (γ) rays) or very small particles (alpha (α) particles and beta (β) particles) emitted from the nucleus of a radioisotope as it decays particles have mass; rays are part of the electromagnetic spectrum of radiation and have no mass alpha, beta, and gamma radiation can be separated using a magnetic field alpha and beta particles have contrary charges and so undergo aberration in opposite directions; gamma rays don't transfer any charge so they don’t undergo aberration Radiation PROTECTION -​ cells are damaged by radiation -​ alpha particles have mass and can not travel into the skin -​ beta particles have very small mass (mass of an electron) and can penetrate 4- 5 nm into body tissue – having clothing acts as a barrier gamma rays pass through body tissues – lead and other heavy metals act as barriers to gamma rays RADON -​ colourless, odourless, radioactive noble gas -​ produced when naturally occurring uranium undergoes radioactive decay -​ can collect in confined areas of the home and contribute to our daily dose of radiation -​ this radiation increases a person’s risk for lung cancer -​ home inspectors test homes for radon gas 1.5 Dobereiner’s Law of Triads (1829) noted similarity among physical and chemical properties of several elements groups of 3 elements noted that some properties of the elements followed trends in which the value of the middle element of the triad was in between those of the end members modern periodic table has these triads of elements in adjacent spaces (either consecutive periods (rows) in a group (column) or consecutive groups in a period) Newlands’s Law of Octaves (1898) arranged elements according to increasing atomic mass periods were shown going down the table, with groups going across – the opposite from the modern form of the periodic table noticed that similar physical and chemical properties appeared for every eight element e.g. Li, Na, K are all soft, grey, highly reactive metals Mendeleev’s Periodic Table (1869) listed all known elements (about 60 at the time) in horizontal rows according to atomic mass, A elements with the same properties appear in the same column (periodic trends) contained blank spots for yet undiscovered elements Van den Broek’s Modern Periodic Table (1911) rearranged elements according to their atomic numbers, Z elements are organized in order of increasing atomic number, yet spaced so that elements with similar physical and chemical properties are aligned in columns PERIODIC LAW physical properties regularly repeat if the elements are arranged in order of increasing atomic number, Z (not atomic mass, A) PERIOD 7 rows, period number = number of occupied energy levels trends in a period are called periodic trends GROUP/FAMILY 18 columns, group number = number of valence electrons similar physical and chemical properties Periodic Table and Electronic Arrangement position of an element is related to electron arrangement in its atom e.g. P: 3rd Period so it has electron in first 3 energy levels (n = 3); and 5th Group so it has 5 valence electrons; electronic structure is 2, 8, 5 LEWIS SYMBOLS a representation of an element consisting of the chemical symbol and dots to represent the valence electrons; electron dot diagram 1.7 Periodic Trends NUCLEAR CHARGE -​ given by atomic number -​ increases by one between successive elements in periodic table as proton (+1) is added to nucleus -​ direct linear relationship EFFECTIVE NUCLEAR CHARGE -​ net force experienced by outer electrons due to positively charged nucleus -​ is less than full nuclear charge SHIELDING outer electrons do not experience full attraction of nucleus because they are shielded from its force by the inner electrons reduces attraction of nucleus for outer electrons reduces effective nuclear charge ELECTRON-ELECTRON REPULSION electrons repel other nearby electrons electrons arrange themselves as far apart as possible to minimize electron-electron repulsion Trends can be explained by: understanding atomic structure... 1. valence electrons 2. nuclear charge (p+): as increase #p+ > increase attraction of e– by p+ 3. shielding (by inner electrons): as increase shielding > decrease ENC 4. electron-electron repulsion: as increase # of e– > increase repulsion bw e– Atomic Radius -​ measurement of the size of an atom -​ distance from center of atom to the outer edge of its electron cloud -​ usually expressed in picometers (pm) (1 pm = 10–12 m);incredibly small! -​ boundary of electron cloud of atom is not defined thus radius is determined indirectly – distance between two nuclei of the same element divided by 2 Trends in AR Down a Group: increases -​ as then number of electrons increases, the number of occupied energy levels (given by period number) increases -​ the inner electrons shield the outer electrons from the positive nuclear force, decreasing the effective nuclear charge, and thus decreasing the attraction of the outer electrons for the nucleus -​ as number of electrons increases, electron-electron repulsion also increases Across a Period: decreases -​ the number of energy levels stays the same while the number of protons in each atom increases (increased effective nuclear charge) and therefore the pull on the valence electrons increases, resulting in a smaller electron cloud Ionic Radius -​ distance from the center of an ion (formed when an atom loses or gains electrons) to the outer edge of its electron cloud -​ positive ion radius: – increases down a group – decreases across a period -​ negative ion radius: – increases down a group – decreases across a period CATION positive ion due to loss of electrons (metals) a positive ion is smaller than its atom – it has *1 less shell* – electron-electron repulsion decreases as remove electrons – increased positive nuclear charge acting on remaining electrons ANION a negative ion due to gain of electrons a negative ion is larger than its atom – *increased electron-electron repulsion* as add more electrons – decreased positive nuclear charge acting on remaining electrons in atom Ionization Energy Ie -​ amount of energy required to remove an electron electrons from an atom or ion in the gaseous state to form a positive ion A(g) + energy > A+(g) + e– -​ a measure of attraction between nucleus and outer electrons -​ can be measured directly and is a property of gaseous atoms FIRST IONIZATION ENERGY, IE1 – removal of 1st electron – first e– to be removed is held much more loosely than last e– - the farther the electrons are from the nucleus, the easier they are to remove Trends in Ie Across a Period: increases – as atomic radius decreases, electrons are held more strongly by nucleus due to increased effective nuclear charge Down a Group: decreases – as atomic radius increases, outermost electrons are farther from nucleus and are shielded by inner electrons, and thus held less strongly IE trend is reverse r trend Electron Affinity -​ the energy released (or absorbed) as a neutral atom gains an electron -​ (if energy is needed to add an electron to a neutral atom, the resulting negatively charged ion will be unstable and will soon lose the electron) -​ if energy is released when an electron is added to a neutral atom, the resulting negatively charged ion will be stable -​ when an electron is added it is attracted by the nucleus and repulsed by the other electrons in the atom -​ energy is released when the attraction on the electron from the nucleus is greater than the repulsion on the electron from the other electrons -​ most first electron affinities are positive (release energy) A(g) + e– > A–(g) + energy Trends in Electron Affinity Across a Period: increases – as radius decreases, attractive force between nucleus and new electron is increased Down a Group: decreases – as atomic radius increases, p+ and e– are farther away and attractive force between nucleus and new electron is decreased – increased shielding by inner electrons Electronegativity -​ relative ability of an atom to attract shared/bonding electron pairs in a covalent bond -​ a relative quantitative value established by Linus Pauling -​ combines ionization energies, electron affinity, bond energy data, and some other measures of reactivity -​ no unit bc a relative measure fluorine is the most electronegative element with EN = 4.0; least is Cs (EN= 0.7) Trends in Electronegativity Across a Period: increases – as nuclear charge increases, atomic radius decreases, and distance between nucleus and electrons decreases, thus attraction between p+ and bond e– increases (e– held more strongly) Down a Group: decreases – as atomic radius increases, distance between nucleus and bond electrons increases, thus attraction between p+ and e– decreases; more shielding (e– held less strongly) NOTE: I vs EN – Although trends in ionization energy and electronegativity are the same, they are distinct properties. Ionization energies can be measured directly and are a property of gaseous atoms. Electronegativity is a property of an atom in a molecule and values are derived indirectly from experimental Bond energy data.