Atoms & Elements

2 Chapter Outline Development of Atomic Theories The composition of Atom Chemical properties of Atoms, the Periodicity and Periodic Table Isotopes

3 Experiencing Atoms Atoms: incredibly small, yet compose everything atoms are the pieces of Elements properties of the atoms determine the properties of the elements

4 Experiencing Atoms 91 elements found in nature Over 20 we have made in laboratories, and scientists are going to make more Each element has its own, unique kind of atoms different structures different properties

5 Dalton’s Atomic Theory 1.Elements are composed of atoms 2.All atoms of an element are identical 3.Atoms combine in simple, whole- number ratios to form molecules of compounds CO 2 (O=C=O): 2 Oxygen atom + 1 Carbon atom H 2 O (H-O-H): 2 Hydrogen atom + 1 Oxygen atom 4.In chemical reactions, atoms are not broken or changed into another type 2H 2 + O 2  2H 2 O John Dalton ( )

6 The Size of Atoms Atomic Mass Unit (amu): 1 amu = 1.66  g Hydrogen the smallest atom mass of H atom= 1.67 x g ~ 1 amu volume of H atom = 2.1 x cm 3

7 About Charge Rubbing dry paper towel and dry plastic piece causes plastic attracting small objects like styrofoam. Benjamin Franklin: Two Kinds of Charge: + and – Opposite Charges Attract: + attracted to – Like Charges Repel + repels + – repels – Neutral: no charge or equal amounts of opposite charges

8 Lightning: Neutralization of Charges accumulated between the clouds and the ground

9 1900s: The Atom is Divisible! Discovery of Electrons (J.J. Thomson et al.): The atom had pieces called electrons -Electrons are much smaller than atoms and carry a “-” charge the mass of the electron is 1/1836 th the mass of a hydrogen atom the charge on the electron is the fundamental unit of charge which we will call –1 charge units

10 What is Inside an atom? (7-10 min) Ernest Rutherford’s experiment (1909): bombardment of a sheet of large atoms (as target) with small, high energy particles bullet = alpha particles, target atoms = gold foil  particles have a mass of 4 amu & charge of +2 c.u. gold has a mass of 197 amu & is very malleable

11 Rutherford’s Experiment Lead Box Radioactive Sample Gold Foil Fluorescent Screen Alpha Particles Striking Screen >98% ~2% 0.01%

12 Rutherford’s Results Thompson’s model of atom predicts there is no heavy mass or high charge within the atom,  particles should penetrate without obstruction > 98% of the  particles: went straight through ~2% of the  particles went through but were deflected by large angles ~0.01% of the  particles bounced off the gold foil: “...as if you fired a 15” canon shell at a piece of tissue paper and it came back and hit you.”

13 Rutherford’s Model of Atoms: Nuclear Atom Atom mostly empty space because almost all the particles went straight through Atom contains a dense particle: small in volume compared to the atom but large in mass because of the few particles that bounced back “+” charge because of the large deflections of some of the particles...

14... Rutherford’s Nuclear Atom Model Thompson’s Plum Pudding Atom was disproved by Experiment Plum pudding model: all the  particles should go straight through most  particles go straight through some  particles go through, but are deflected a few of the  particles do not go through

15 Modern View of Atom Atoms = (Protons + Neutrons) + Electrons The nucleus: Protons + Neutrons The nucleus is only about cm in diameter The electrons move outside the nucleus with an average distance of about cm therefore the radius of the atom is about 100,000 times larger than the radius of the nucleus

16 Inside the Nucleus: Neutron and Protons Protons: “+” charge and a mass of 1 amu the number of proton equals the number of electrons in an atom (electrically neutral) Neutron: have no charge and a mass of 1 amu the masses of the proton and neutron are both approximately 1 amu

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18 Elements each Element has a unique number of protons in its nucleus Atomic number: the number of Protons in the nucleus of an atom the elements are arranged on the Periodic Table in order of their atomic numbers each element has a unique name and symbol symbol either one or two letters  one capital letter or one capital letter + one lower case

19 The Periodic Table of Elements

20 Get to know Elements: Name, Symbol, Atomic Number Element NameSymbolAtomic Number Iodine Sb 82

21 Did you find them? What is the atomic number of boron, B? ____ What is the atomic mass of silicon, Si? ____ How many protons does a chlorine atom have? ____ How many electrons does a neon atom (neutral) have? ____ Will a Na atom with 10 electrons be electrically neutral? ____

22 Mendeleev: Periodicity order elements by atomic mass  repeating pattern of properties Periodic Law – When the elements are arranged in order of increasing relative mass, certain sets of properties recur periodically used pattern to predict properties of undiscovered elements A good documentary: Mendeleev (in the first 25 minutes, min as my fav )

23 An example of Periodic Pattern: Compound with oxygen (E x O y ) or with hydrogen (EH x ) H nm H 2 O a/b 1 H 2 Li m Li 2 O b 7 LiH Na m Na 2 O b 23 NaH Be m/nm BeO a/b 9 BeH 2 m MgO b 24 MgH 2 Mg nm B 2 O 3 a 11 (BH 3 ) n B m Al 2 O 3 a/b 27 (AlH 3 ) Al nm CO 2 a 12 CH 4 C nm/m SiO 2 a 28 SiH 4 Si nm N 2 O 5 a 14 NH 3 N nm P 4 O 10 a 31 PH 3 P nm O 2 16 H 2 O O nm SO 3 a 32 H 2 S S nm Cl 2 O 7 a 35.5 HCl Cl nm OF 2 19 HF F metal (m) or nonmetal (nm) oxide: acidic (a) or basic (b) Compound w/ H

Periodicity = Metal = Metalloid = Nonmetal

25 Metals: Physical vs. Chemical Properties solids at room temperature, except Hg reflective surface shiny conduct heat, electricity Malleable (can be shaped) Tend to Lose electrons and form Cations in reactions. Na  Na + + e - about 75% of the elements are metals lower left on the table

26 Nonmetals : Physical vs. Chemical Properties Elements found in all 3 states poor conductors of heat or electricity solids are brittle Tend to gain electrons in reactions to become anions: Cl + e -  Cl - upper right on the table except H AtRn IXe SeBrKr PSClAr CNOFNe He

27 Metalloids: between Metals and Nonmetals show some properties of metals and some of nonmetals also known as semiconductors Properties of Silicon shiny conducts electricity does not conduct heat well brittle Po SbTe GeAs Si B IIB IIIA IVA VA VIA VIIA VIIIA

28 The Modern Periodic Table Elements with similar chemical and physical properties are in the same column columns are called Groups or Families designated by a number and letter at top rows are called Periods each period shows the pattern of properties repeated in the next period

29 The Modern Periodic Table Main Group = Representative Elements = ‘A’ groups Transition Elements = ‘B’ groups all metals Bottom rows = Inner Transition Elements = Rare Earth Elements metals really belong in Period 6 & 7

30 = Alkali Metals = Alkaline Earth Metals = Noble Gases = Halogens = Lanthanides = Actinides = Transition Metals

31 = Rare Earth Metals = Transuranium element = Transition Metals U

32 Important Element - Hydrogen nonmetal colorless, diatomic gas H 2 very low melting point & density reacts with Nonmetals to form molecular compounds HCl is acidic gas H 2 O is a liquid reacts with Metals to form hydrides metal hydrides react with water to form H 2 Nickel-metal hydride (NiMH) used in rechargeable battery HX dissolves in water to form acids

33 Important Groups – IA, Alkali Metals Usually Hydrogen is included All metals: soft, low melting points Flame tests  Li = red, Na = yellow, K = violet Chemical Property: Very reactive. React with water to form basic (alkaline) solutions and H 2.React with water to form basic (alkaline) solutions and H 2.  releases a lot of heat Tend to form water soluble compounds, such as table salt and baking soda.  colorless solutions lithium sodium potassium rubidium cesium

34 Important Groups – IIA, Alkali Earth Metals Physical properties: harder, higher melting, and denser than alkali metals flame tests  Ca = red, Sr = red, Ba = yellow-green Chemical properties: reactive, but less than corresponding alkali metal form stable, insoluble oxides. oxides are basic = alkaline earth reactivity with water to form H 2,  Be = none; Mg = steam; Ca, Sr, Ba = cold water magnesium calcium beryllium strontium barium

35 Important Groups – VIIA, Halogens nonmetals F 2 & Cl 2 gases; Br 2 liquid; I 2 solid all diatomic very reactive Cl 2, Br 2 react slowly with water Cl 2 + H 2 O  HCl + HOCl (chlorine) react with metals to form ionic compounds HX all acids HF weak < HCl < HBr < HI bromine iodine chlorine fluorine

36 Important Groups – VIIIA, Noble Gases all gases at room temperature, very low melting and boiling points very unreactive, practically inert very hard to remove electron from or give an electron to

37 Ion: Charged Atom The number of protons determines the element! all sodium atoms have 11 protons in the nucleus In a chemical change, the number of protons in the nucleus of the atom doesn’t change! no transmutation during a chemical change!! during radioactive and nuclear changes, atoms do transmute Atoms in a compound are often electrically charged, these are called ions

38 Ions Atoms acquire a charge by gaining or losing electrons not protons!! Ion Charge = # protons – # electrons ions with a + charge are called cations more protons than electrons form by losing electrons ions with a – charge are called anions more electrons than protons form by gaining electrons

39 Atomic Structures of Ions Metals form cations More positive charge, the fewer electrons than the neutral atom Na atom = 11 p + and 11 e -, Na + ion = 11 p + and 10 e - Ca atom = 20 p + and 20 e -, Ca 2+ ion = 20 p + and 18 e - Cations are named the same as the metal sodiumNa  Na + + 1e - sodium ion calciumCa  Ca e - calcium ion The charge on a cation can be determined from the Group number on the Periodic Table Group 1A  +1, Group 2A  +2, (Al, Ga, In)  +3

40 Atomic Structures of Ions Nonmetals form anions For each negative charge the ion has 1 more electron than the neutral atom F = 9 e -, F - = 10 e - P = 15 e -, P 3- = 18 e - Anions are named by changing the ending of the name to -ide fluorineF + 1e -  F - fluoride ion oxygenO + 2e -  O 2- oxide ion The charge on an anion can be determined from the Group number on the Periodic Table Group 7A  -1, Group 6A  -2

41 Example: Find the number of protons and electrons in the Ca 2+ ion number of protons = _____ Number of electrons = _____ The atomic number of Calcium is the number of ________ons. The “2+” charge means the ion has two _______ electrons The number of electrons in an ATOM equals the number of _____ons. Therefore the number of electron in Ca 2+ ion = ____ __ ____ = ___

42 Example: Find the number of protons and electrons in an I - ion number of protons = _____ Number of electrons = _____ The atomic number of iodine is the number of ________ons. The “-” charge means the ion has ONE _______ electrons The number of electrons in an ATOM equals the number of _____ons. Therefore the number of electron in I - ion = ____ __ ____ = ___

43 Atomic Structures of Ions

44 Ion Charge & the Periodic Table the charge on an ion: an elements position on the Periodic Table Metals: always positive ions Na + Ca 2+ Al 3+ Nonmetals are negative ions Cl - O 2- N 3- Main group metals: #charge = #group Nonmetals, #charge = #group - 8

45 Li + Na + K+K+ Rb + Cs + Be 2+ Mg 2+ Ca 2+ Sr 2+ Ba 2+ Al 3+ Ga 3+ In 3+ O 2- S 2- Se 2- Te 2- F-F- Cl - Br - I-I- N 3- P 3- As 3- IA IIAIIIA VIIA VIA VA Charge of ions by Group __ __ __

46 Same Element, Different #Neutrons? Isotopes: The same element could have atoms with different masses 2 isotopes of chlorine found in nature: one has a mass of about 35 amu (Cl-35); another that weighs about 37 amu (Cl-37) Carbon-12 is much more abundant than Carbon- 13. C-14 formed from nuclear reaction of N-14. The observed mass is a weighted average of the weights of all the naturally occurring atoms the atomic mass of chlorine is amu

47 Isotopes all isotopes of an element: chemically identical undergo the exact same chemical reactions the same number of protons different masses due to different numbers of neutrons. Example: C-14 atom has eight neutrons; C-12 atom has six neutrons. identified by their mass numbers protons + neutrons

48 Atomic Number (Z) Number of protons Mass Number (A) Protons + Neutrons Abundance = relative amount found in a sample Example: Cl-35 (75%) vs. Cl-37 (25%) Isotopes

49 Isotopic Symbol Cl-35 has a mass number = 35, 17 protons and 18 neutrons ( ). The symbol for this isotope would be Atomic Symbol A = mass number Z = atomic number #neutrons = A - Z AXAX Z Cl 35 17

50 Write the Isotopic symbol for Cr isotope w/ 27 neutrons #protons = _____ #mass = _____ Write down the given quantity, questions, and equations Given: Element, #neutron (#n) Find: #proton (#p) Find: #mass (A) Given: #p = Z, A = #p + #n

51 Example: How many protons and neutrons in the chromium isotope #proton = __ #neutron = _________ = ____

52 Practice - Complete the following table

53 Mass Number is Not the Same as Atomic Mass the atomic mass is an experimental number determined from all naturally occurring isotopes the mass number refers to the number of protons + neutrons in one isotope natural or man-made

55 Atomic Mass is the Weighted average of Mass Number Natural abundance of isotope: The percentage of #atoms for an isotope among all the naturally existing isotopes. Gallium: Ga-69 (mass number amu, abundance 60.11%); Ga-71 (mass number , abundance 39.89%). Atomic mass of Gallium = amu x x = amu