Chapter 8 - Atoms and Periodic Properties

1. Chapter 8 - Atoms and Periodic Properties Will turn to a study of the properties of matter why materials have certain properties chemistry - composition, structure and properties of substances and the transformations they undergo consider world - many objects with many properties trees bark, leaves, wood car wheels, dash, hood all substances made of combinations only about 113 known elements element - pure substance that cannot be decomposed into simpler substances by a chemical or physical process well-defined properties Water - H2O Salt - NaCl

2. + Periodic Table - lists all known elements currently about 108 known elements 88 naturally occurring others made in lab Some substances known for long time how to describe? alchemists - “lead into gold” antimony Sb confusing STANDARDIZATION - modern symbols used world-wide now how did we get these symbols? H He C Cl B Be O Os P Pt S Se N Ni Hg Ag Au Na Fe Sn Pb

3. Where do the names come from? Pu Es Am U Fm Fr Hg Md Eu Np Bh Cf Cl - light green Tc - artificial Ne - new He - sun Te - Earth Elements named after planets, people, places and descriptions! Names passed by international council “Commission on the Nomenclature of Inorganic Chemistry” Names agreed upon worldwide standardized

4. + + + + + - - - - - Elements made up of large collection of atoms Atom - smallest unit with the same chemical identity as element (10-10 m, 10-24 g) chemical identity - physical and chemical properties of a pure substance structure of the atom: protons, neutrons, and electrons NUCLEUS - fixed central part of atom contains: proton positive charge strong nuclear force neutron no charge same mass as proton does not affect chemical identity these are also called nucleons-reside in nucleus electron - negative charge (same as proton) -swarm around nucleus (electron cloud) -can be attracted away or added w/o chemical change -very light 1/1837 mass of proton (negligible) neutral atom-same number of p+ as e- (zero net charge) remember ion: atom with net charge

5. How to determine atomic structure History: Ancient Greeks Democritus-matter is discontinuous cannot divide indefinitely “atom” - Greek for uncuttable Aristotle & Plato disagreed with this view (wrongly) thought matter was continuous John Dalton (1800’s) revisited idea of Atoms Dalton’s Atomic Theory • All matter = indivisible atoms • An element is made up of identical atoms • Different elements have atoms with different masses • Chemical compounds are made of atoms in specific integer ratios • Atoms are neither created nor destroyed in chemical reactions

6. MODERN IDEAS – discoveries leading to atomic structure – indirect observations J.J Thompson (late 1800’s) – discovery of electrons cathode ray tubes – eject particles from plates -cathode rays found to be negative (opposites attract-not light) -deflect in magnetic field (current-moving charge) - measured charge-to-mass ratio (crossed electric&magnetic fields) Robert Millikan (1906) Oil drop experiment -charged oil drops in electric field -electric force opposed gravity – drop floats -droplet charge in multiples of electron chargeqe=1.6x10-19 C -found electron mass by using q/m from Millikan me=9.11x10-31 kg very very small

7. Early model of the atom Plum pudding model Electrons embedded in blob of positively charged matter like “raisins in plum pudding” But what is the positive charge that cancels tiny electrons? Rutherford – alpha particle positive helium nucleus scattering- shoot alpha particle at gold sheet Result : -most of the alpha particles passed through sheet -some alpha particles back-scattered Conclusion: -atom contains small central part most of mass nucleus -electrons orbit at distance 100,000 times the size of the nucleus the atom is mostly made up of EMPTY SPACE Nucleus later found to be made of protons (Rutherford split nucleus) and neutrons (Chadwick-1932)

8. + + + + + + + + + Describing the Modern Atom atomic number – number of protons in nucleus -describes identity of element -neutral atom number of e- = number of p+ mass number – number of protons and neutrons in nucleus indicates mass since the electrons are negligible new mass scale – STANDARDIZE atomic mass units (amu, dalton) 1 amu is about mass of a proton amu defined by mass of carbon-12: carbon 12: 6 protons and 6 neutrons define to have mass of exactly 12 amu ATOMIC MASS STANDARD But mass number does not define element can have different numbers of neutrons For example: Lithium Li ATOMIC NUMBER: 3 3 3 MASS NUMBER: 3 amu 4 amu 5 amu Isotope: elements with the same number of protons, but different numbers of neutron – different mass numbers

9. How do we study isotopes? Mass Spectrometer Curve of ions depends on the charge-to-mass ratio -isotopes have different masses N S Each isotope will form a spot At different places on the screen Ions accelerated in electric field oven Natural Abundance-what percentage of each isotope exists for each element Mass number –refers a particular isotope - specific atoms Atomic Weight (Mass) – weighted average of the masses for different isotopes in a sample of an element for the element in general (all isotopes) Notation: describes atomic structure: for an isotope number of protons, neutrons and electrons mass number Means atomic number=78 and mass number =112 amu 11278Pt Atomic structure- protons: 78 electrons: 78 neutral neutrons: 112-78=34 atomic number Example: 3517Cl

10. An important isotope : H lightest element H D T hydrogen deuterium tritium 11H 21H 31H normal heavy radioactive hydrogen hydrogen hydrogen-1 hydrogen-2 hydrogen-3 Natural 99.98% 0.015% < 0.005% Abundance atomic weight: 1.008 amu Remember atomic model: SOLAR SYSTEM MODEL massive nucleus surrounded by electrons problem: electron circles atom - centripetal acceleration classical charge radiates if accelerated loses energye-falls into nucleus New Theory needed -- F=ma didn’t work Planck & Einstein : matter absorbs discrete amounts of energy QUANTA

11. BOHR MODEL: tried to match experiments involving absorption and emmision of light from hot solids and gases - line spectra not derived but phenomenological Bohr’s Theory: 1. Electrons orbit the nucleus at specific distances from the nucleus -allowed orbits 2. Electrons in allowed orbits do not radiate energy -contrary to classical theory 3. Electrons gain energy by “jumping” to a higher energy (further) orbit -lose energy by falling to a lower energy -energy loss or gain in the form of a photon- particle of light “Qnantum Leap” n=1 n=2 n=3 Explained line spectra - electrons in matter gain (absorb) or lose (emit) photons to make quantum leaps

12. Wave-Particle duality : light travels like particles and waves de Broglie : matter also travels like waves electrons travel like waves -normal objects have very small wavelength -electron motion governed by wave properties STANDING WAVE SOLUTION Only certain wavelengths (energies) will fit correctly around nucleus ALLOWED ORBITS Led to the development of QUANTUM MECHANIC THEORY Schrodinger Equation-solve with linear algebra and differential equations Solution: electron orbital - 3D region surrounding nucleus where there is the greatest probability of finding an electron

13. Consequences of quantum mechanics Solution gives energy levels of electrons surrounding nucleus -gives electron configuration -the arrangement of electrons in orbitals and suborbitals about the nucleus of an atom -describes properties of atom “fingers of the atom” interact through electrons PROBABILITY DENSITY - probality at a particular position Cannot isolate position of an electron HEISENBERG UNCERTAINTY PRINCIPLE cannot measure momentum (motion) and position of electron exactly SOLUTION TO WAVE EQUATION GIVES QUANTUM NUMBERS -describe energies of the electrons -determine properties of electrons in atom -gives framework to “build” atoms -similar electron configuration gives similar properties -restrictions on what numbers can be

14. n=1 n=2 n=3 QUANTUM NUMBERS - n,m,l,s describe energy levels of electrons Principal Quantum Number ( n ) main energy level of electron -describes orbit -how far electron is from nucleus -similar to allowed Bohr orbit -restriction: whole number greater than 0 n=1,2,3,4,… closest furthest Angular Momentum Quantum Number ( l ) shape of electron orbit -how spread out the orbital is -restriction:l= 0 to (n-1) l = 1 hourglass l = 0 sphere

15. Magnetic Quantum Number ( m ) orientation of electron orbital -the way the electrons are oriented about nucleus -restriction:-l > m > +l Example : n=2 electrons l=0 electrons have possible m=0 only one way to orient sphere l=1 electrons have possible m= -1, 0, +1 oriented in y-dir m=-1 oriented in x-dir m=+1 oriented in z-dir m=0 n, l, m describe spatial properties of electron how the electron cloud looks each quantum number describes electron with specific energy in nucleus!

16. - - - - - - - - Spin Quantun Number (s) magnetic properties of the electron electron - electric charge spin clockwise spin counterclockwise spin up spin down electron magnets interact with magnetic field split into two beams in magnet - restriction s = +1/.2, -1/2 ELECTRON CONFIGURATION determined by the values of quantum numbers n, l, m, s “fingers of the atom” how the electrons interact with their environment PHYSICAL AND CHEMICAL PROPERTIES (CHEMICAL IDENTITY) S N N S

17. CAN NOW BUILD ATOMS Need some rules before building electron configurations Electrically neutral atoms - same number of e- as p+ have no net charge Ground state atom - electrons occupy only the lowest energy levels in atom as opposed to excited state - electrons occupy higher energy states lower energies unoccupied Will add electrons up to atomic number -but how do we add electrons? TWO WAYS TO ADD PARTICLES: 1. Put all electrons in lowest energy level {n=1, l=0, m=0, s=1/2} cannot do for e- only for BOSONS (photons) 2. Pauli Exclusion Principle: no two electrons in the same atom can have the same set of quantum numbers n, l, m, s Law of nature for fermions (spin=1/2)

18. BUILD TABLE OF ELEMENTS First note: for the spatial orbital n, l, m describe position in space two values of s for each nlm combination Start building - add electrons successively to each lowest energy orbital H : atomic number = 1 one electron put electron in lowest energy Electron configuration: n=1, l=0, m=0, s=1/2 He : atomic number = 2 two electrons Electron configuration: n=1, l=0, m=0, s=1/2 still lowest n=1, l=0, m=0, s=-1/2 -different atom just change to s=-1/2 (next energy) LI : atomic number = 3 three electrons Electron configuration: n=1, l=0, m=0, s=1/2 n=1, l=0, m=0, s=-1/2 n=1 full, next n=2 n=2, l=0, m=0, s=1/2 electron capacity Be to Ne are filled by adding two e- to each n, l, m Electron capacity - maximum number of electrons that can be added to each orbital

19. Connections to Periodic Table Note: row (period) determined by highest principal quantum number electron capacity met at end of row: NOBLE GASES outer shell not full - chemically reactive outer shell full - no electrons to interact with other elements chemically inert NOBLE GAS INERT GAS get to next element by adding electron to next level in orbital up to electron capacity Electron properties determined by principle (n) and angular momentum (l) Q.N Electron orbital notation: specify n, l, and number of electrons in the orbital (superscript) electron capacity s orbital l =0 2 p orbital l =1 6 electrons d orbital l =2 10 allowed in the f orbital l =3 14 orbital Example : 3d2 n=3, l =2, 2 e- in orbital

20. Rewrite electron configuration in new notation H Li Na K Be Mg Ca B Al C Si N P O S F Cl He Ne Ar Sc next: things are different, but first Periodic (Moseley’s) Law - electron configurations repeat properties of elements repeat when ordered by increasing atomic number periodic function of atomic number similar outer shell, similar properties First column : Alkali metals s1 orbital very reactive - single electron Second column: Alkaline Earth Metals s2 orbital Last column : (Inert) Noble Gases s2p6 (s2 for He) very stable - octet (eight outer e- except He) Electron configuration repeats, chemical properties repeat

21. Led to Periodic Table Mendelev thought to be the father of the Periodic Table Periodic Table – table of all know elements listed in order of atomic number -periodicity in properties along rows (density, melting/boiling, hardness, etc) DIVIDED INTO: Families (Groups) – vertical column of elements these elements exhibit similar properties have same outer electron configuration Eight main groups : Group IA to Group VIIIA MAIN GROUP or REPRESENTATIVE GROUPS show similarities in outer e- shell ( want octet ) Group IA – Alkali Metals (react violently w/H2O) s1 never uncombined in nature Group IIA – Alkaline Earth Metals (also reactive) s2 Group IIIA s2p1 Group IVA s2p2 Group VA s2p3 Group VIA s2p4 Group VIIA –Halogens (salt former w/metal) s2p5 Group VIIIA - Noble (Rare)Gasses s2p6 never bond with others { explains Periodic Law

22. Transition Metals – B groups Group IB to Group VIIB fill inner electron orbitals like Sc: 1s22s22p63s23p64s23d1 skips energy change in orbital energies higher orbitals have lower energy i.e., 4s is lower than 3d shows gap in periodic table electron energy level order: 1s2s2p3s3p4s3d4p5s4d5p6s4f5d6p7s5f6d7p6f7d increasing energy can write any electron configuration Example 1. Write configuration for: Zr V Example 2. Identify element electron configuration 1s22s22p63s23p64s23d7 1s22s22p63s23p4 Periods – row groupings of the Periodic Table properties repeat as you go from one period to the next -periods begin reactive (IA) and end stable (VIIIA) Note: Period and Group of element identified properties Historically: some elements undiscovered- chemists knew properties before it even existed

23. How to read Periodic Table Name Group II (family): 2 e- Atomic number group: # of outer e- Symbol Atomic weight Period 4 Magnesium 12 Mg 24.31 Electron dot notation KERNEL – nucleus and inner electrons -dots represent the outermost electrons -shows what’s available for the other atoms to interact with C group IV 4 outer electrons Metals non-metals and semiconductors (semimetals) METALS- conducts heat and electricity - metallic luster (shiny) - maleable pond into sheets - ductile  draw into wires (extrusion) -form positive ions by losing electrons Li 1s22s1 Li+ stable [He] config Mg 1s22s22p63s2  Mg+2stable [Ne] octet NONMETALS- insulators - dull appearance - brittle - form negative ions to complete octet Cl 1s22s22p63s23p5+1 Cl-stable [Ar] config X X