atoms, molecules and ions. i. atoms & molecules a. atom examples - matter around us is made up...

Atoms, Molecules and Ions

I. Atoms & Molecules A. Atom Examples

- Matter around us is made up of tiny particles called atoms. There are ~ 118 known atoms and these are given in the Periodic Chart.

- The atoms have names and they have one or two letter symbols.

- The following is a Game to help learn a few symbols and placement on the Periodic Table - the answers to this quiz are the names of the correct elements.

Half a dime = Lone Ranger’s Horse =

What I do when I am hungry = A frivolous prisoner =

Storage place for streetcars =

I. Atoms & Molecules A. Atom Examples Continued

- A male member of the Ganese tribe = (Transition

Metal)

- What to do with an ailing man – two answers = (Columns 18&2)

- Comment made at end of box of candy: “oops, they ..” (Col 18)

- What some science classes do, but not this one = (Col 13, Row 2)

- Atoms can exist as:

1) Neutral elements: He Na Cu Hg

2) Negative ions (anions): F-1 or F- I- O-2

3) Positive ions (cations): Na+1 or Na+ H+ Al+3

-

I. Atoms & Molecules B. Molecules

- Atoms combine together in definite ratios to form new units called Molecules (Molecular Compounds) & Formula Units (Ionic Compounds) - We represent Molecules & Formula Units with the symbols of the elements (left to right from Periodic Table) and with sub numbers to indicate the number of atoms other than one.

- General Example: # ZxTy Where # equals the number of molecules and x & y equal the number of atoms in one molecule. Generally place element closest to group I first [except for organic molecules with C first, H second, rest in alphabetical order]. If a compound has more than 1 polyatomic ion, then use parenthesis.

- Few Examples: NaF CO2 H2O C2H4Br2

Ca3(PO4)2 3 H2SO4 C256H381N65O79S6 (Insulin)

I. Atoms & Molecules C. Size

H+ ~ 10-13 cm diam. Ho ~ 10-8 cm diam. Mass H = 1.67x10-24 g

Note: 6.02x1023 = Avogadro’s Number (6.02x1023)x(1.67x10-24g)=1.00 g

- The following are images of 1) Electron Microscope

2) Latex Molecules 3) C in Graphene 4) U atoms 5) Fe on Cu

Atoms & Molecules D. History

~400 BC - Democritus suggested the existence of atoms.

66 AD - Peter wrote: “but the day of the Lord will come like a thief, in which the heavens will pass away with a roar and the elements will be destroyed with intense heat, and the earth and its works will be burned up.” From New Testament, 2 Peter 3:10.

1783 - Antoine Lavoisier found that matter is neither created nor destroyed in a chemical reaction. Known as “father of modern chemistry.”

Atoms &Molecules D. History

1803 - John Dalton proposed that matter is made up of tiny atoms; that atoms of the same element are alike; & that atoms combine in definite ratios to form compounds. This set aside false idea promoted by Aristotle 2000 years earlier that matter was continuous, and reaffirmed Democritus’s early “atomic model.”

I. Atoms & Molecules D. History

1879 - William Crookes developed the “ray tube” which later allowed us to view electron beams


1897 – Joseph Thomson used the cathode-ray tube and discovered the electron.


1886 - Eugene Goldstein demonstrated existence of + particles, protons. These particles later found to have a charge of +1 (1.60x10-19 coulombs) and a mass of 1.67x10-24 g (a mass of 1.00 AMU).

1909 - Robert Millikan determined mass (9.11x10-28 g; ~1800 less than proton) and -1 charge (-1.60x10-19 coulombs) of an electron.


1911 - Ernest Rutherford (a New Zealand physicist) demonstrated the nuclear nature of the atom in which the empty space is 10,000 to 100,000 times larger than the size of the nucleus.


1932 - James Chadwick demonstrated the existence of the neutron which has no charge and about the same mass as the proton (1.00 AMU).

- Why do you think that it took longer to uncover the neutron than either the electron or proton?

I. Atoms & Molecules Summary

Atoms are made up of three major parts:

Part Found Mass Charge

Protons Part of Nucleus 1.7x10-24g (1.0 AMU) +1

Neutrons Part of Nucleus 1.7x10-24g (1.0 AMU) 0

Electrons Outside Nucleus 9.1x10-28g (small) -1

Notes: 1) Neutral atoms contain equal # of electrons and protons.

2) Atoms can loose or gain electrons to become charged = ions

3) Electrons are reshuffled in a chemical reaction.

4) # protons (Atomic #) determines the identity of the atom or ion.

5) Mass of atom = # Protons plus # Neutrons

I. Atoms & Molecules E. Atomic Structure

- Heavier protons and neutrons in the center or nucleus and the smaller electrons are found outside of the nucleus in shells of specific energy.

- The outer electrons and their arrangements in specific energy levels are responsible for the chemistry of the element.

- The number of protons (atomic number) determines the identity of the atom – NOT the number of neutrons; NOT the number of electrons; and NOT the atomic weight.

- The number of protons plus neutrons determines the weight of the atom in AMU. Atomic Mass = #n + #p (from periodic chart). Why are the masses not integers?

- The ratio of electrons to protons determines the charge of the atom; atoms can loose or gain electrons to have + or - charges; for neutral atoms the #e = #p.

I. Atoms & Molecules E. Atomic Structure Continued

- Atoms of a given element with differing numbers of neutrons are called isotopes. For example: ( note Mass #

Atomic # H )

11H = 1P & 0N 2

1H = 1P & 1N (deuterium) 31H = 1P 2N (tritium)

- Note that the weights of the atoms are averages of the masses of the naturally occurring isotopes.

- For example: 52.0 % of Br has 44 neutrons = mass 79.0 [ 7935 Br ]

48.0 % of Br has 46 neutrons = mass 81.0 [ 8135 Br ]

Average mass = (0.520 x 79.0) + (0.480 x 81.0) = 79.9

- Note that the masses of the atoms and their isotopes are determined with a useful instrument called a mass spectrometer ( MS ) .

I. Atoms & Molecules F. Mass Spectrometer – Used as a GC & LC Detector for qualitative & quantitative analysis of atoms and compounds. See Pgs 98 & 99 for information & a mass spectrum of CH2Cl2

+ + +

-

-Negative Plate

Beam of Electrons

+ ions separatedand detected

Mass Spectrometer

Sample Introduced to the MS frequently by GC or LC

Simplified Mass Spectrometer (MS of Ne)

I. Atoms & Molecules F. Mass Spectrometer Continued

Result = mass spectrum; a plot of intensity versus m (m/e).Qual. from masses and Quant. from intensities of peaks.Example: MS of a mixture of NaCl and NaBr; plot of

intensity (y axis) vs mass (x axis).

23Na+ 35Cl+ 37Cl+ 79Br+ 80Br+

Intensity

Mass

Element ppb Element ppb Element ppbLi 1.8 Be 2.3 B 3.0Na >1000 Mg >1000 Al 10Si >1000 P 28 S >1000Cl 70 K >1000 Ca >1000Sc 8.2 46Ti 35 48Ti 73V 32 Cr 21 Mg 33Fe >1000 Co 21 Ni 3863Cu 13 65Cu 17 66Zn 3768Zn 40 Ga 7.6 Ge 2.6As 10 Se 13 Br 18Se 8.9 Rb 1.6 Sr 200Zr 0.6 Nb 0.9 Mo 120Ru 1.3 Rh 0.4 Pd 1.6Ag 13 Cd 27 In 0.7Sn 0.6 Sb 1.1 I 0.5Te 0.4 Ba > 1000 La 0.4Ce 0.2 Pr 0.5 Nd 1.7Sm 0.8 Eu 0.3 Gd 0.9Tb 0.3 Dy 0.6 Ho 0.1Er 0.6 Tm 0.4 Yb 0.7Lu 0.2 Hf 1.0 Re 0.3Os 0.4 Ir 0.5 Pt 0.6Au 0.2 Hg 3.1 Tl 6.5Pb 22 Bi 8.0 Ur 0.4

Poisoned Pop & ICP-MS Results

II. Periodic Table A. Introduction

1869 - Dmitri Mendeleev (Russian chemist) & Julius Lothar Meyer (German chemist) independently arranged the known atoms by atomic weight (now by atomic number).

II. Periodic Table A. Introduction

- Dmitri Mendeleev & Lothar Meyer independently arranged the known atoms by atomic weight in 1869.

- In early 1900’s we rearranged the Periodic Chart by atomic number and found that numerous physical and chemical properties had a regular repetition.

- Rows are called periods; are 7 of these; named 1 to 7; properties change going across a row.

- Columns are called groups; are three names used: (IA--VIIIA, 1--18, and common names); all elements in a group have similar chemical properties.

- Following are some periodic properties.

II. Periodic Table B. Periodic Properties 1. Metals & Nonmetals

- Metals on Left Side; Note dividing line (H = nonmetal); Good conductors of heat & electricity; Poor insulators; Lustrous; Solids; Malleable; Lose Electrons to form Cations like Na+ Mg+2 Fe+3 Al+3

- Nonmetals on right side: Poor conductors of heat & electricity; Not lustrous; Brittle if solid; Many are gases; Can gain electrons to form anions, & can also share electrons to yield neutral molecules:

F- O-2 F2 [F – F] H2 [H – H] O2 [O = O]

B C Nonmetals

Al Si

Metals Ga Ge As H IS A NON-

METAL

In Sn Sb Te

Tl Pb Bi Po At

II. Periodic Table B. Periodic Properties 2. Size

- Neutral atoms get larger as go down periodic chart.

- Neutral atoms get larger as go left on periodic chart.

Fr

II. Periodic Table B. Periodic Properties

3. Ionization Energy

- Energy needed to remove outer electron of a neutral atom.

- Generally increases as go up and as go to right on PC.

4. Electronegativity (Ignore group 18 – helium group)

- A measure of the desire of atom for electrons.

- Gets larger as go up and as go right on the periodic chart;

F is most and Fr is least electronegative. Important

F

Fr


5. Acidity/Basicity: (Will define later)

- Elements on left form bases like NaOH, Ba(OH)2

- Elements on right form acids like HCl, HNO3, H2SO4

6. Charges: Some atoms lose or gain electrons to mimic

nearest group VIIIA/18 element. Know these charges.

IA/1 IIA/2 IIIA/13 VA/15 VIA/16 VIIA/17 VIIIA/18

+1 +2 +3 -3 -2 -1 0


- Note: Under normal conditions, elements in Black are solids (K, S, U); elements in Blue are liquids (Hg & Br2); elements in Red are gases (H2, F2, Ar); outlined elements are man-made (Tc, Pu).

- IMPORTANT Memorize: Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine, Bromine & Iodine normally exist as diatomic molecules, and these molecules have the same names as the elements:

H2 O2 F2 Cl2 Br2 I2

hydrogen, oxygen, fluorine, chlorine,……………….

III. Compounds A. Introduction

- Atoms combine in definite ratios to form compounds.

- Are two classes of compounds:

1) Ionic Compounds (Compound containing a METAL)

– Metals lose electrons in chemical reactions to form + ions called cations. This will occur when the metal reacts with a nonmetal. Nonmetals gain electrons to become - negative ions called anions. (Note: H classified as a nonmetal)

- The cations and anions attract each other in such a ratio so that the compound is neutral.

- Ionic compounds form aggregates and we give the simplest ratio of the atoms, called the formula unit.

- Examples: Na+ + F- -----) NaF (Don’t write as Na+F-)

K+ + K+ + O2- -----) K2O


Predicting formulas for ionic compounds- Need to memorize general charges for the groups:

IA/1 IIA/2 IIIA/13 VA/15 VIA/16 VIIA/17

+1 +2 +3 -3 -2 -1- Some metals have variable charges:

Cu+1 Cu2+ Fe2+ Fe3+ Cr3+ Cr6+ Hg2+ Hg22+

Pb2+ Pb4+ Sn2+ Sn4+ Need to know these.

- Add the anions and cations in the simplest ratio to get a neutral compound. Do not show charges in formula.

- Examples: K and F Be and O Ba and Cl Fe3+ and S Mg and N Al and F

(KF BeO BaCl2 Fe2S3 Mg3N2 AlF3)


2) Molecular (Compound with All NONMETALS)

- H and other nonmetals can chemically bind with other nonmetals by sharing electrons to produce new units called molecules, and the molecular formula gives the exact number of atoms in the molecule.

- Ionic compounds existed as aggregates that we show as formula units; however, nonmetals combine to produce distinct molecules. Molecules are represented with molecular formulas or with structural formulas. Structural formulas show atoms attached with covalent bonds; each bond is a line and represents two shared electrons.

- Examples:

H2O2 or H-O-O-H CO2 or O=C=O


Organic Compounds- One class of molecular compounds is organic compounds.- Carbon containing compounds plus H and frequently O, N & Cl

are called organic compounds. Write C, then H, then rest of elements in alphabetical order: C3H9NO

- The vast majority of known chemicals are organic; there are an infinite number of possible organic compounds.

- Organic compounds are organized by functional group (that portion of the organic compound that governs its chemistry).

Examples:

Ether: R – O – R such as CH3-CH2-O-CH2-CH3 (C4H10O)

Alcohol: R – O – H such as CH3-OH (CH4O)

Amine: R – NH2 such as NH2-CH2-CH2-CH2-CH2-NH2 (C4H12N2)

III. Compounds B. Examples

- Classify the following as either ionic or molecular:

NaF CO2 BaF2 H2O SF2 H2 FeI2 CH4

- Predict the formula when the following atoms & ions form ionic substances:

Na & I Ba & I K & O Ca & S

Be & F Be & Se Na & N Mg & N

Fe+3 & I Fe+2 & I Fe+3 & O Fe+2 & O

III. Compounds C. Polyatomic Ions

We frequently encounter ions in which several atoms have combined together to form a polyatomic ion.

Polyatomic Ion: An ion consisting of several atoms bonded together and carrying a charge.

Some of these polyatomic ions occur so frequently that we need to know the names, formulas and charges on the ions.

When have more than one polyatomic ion, then use ( ).Examples:

NaOH OH-1 = Hydroxide

KNO2 NO2-1 = Nitrite

Ca(NO3)2 NO3-1 = Nitrate

III. Compounds C. Polyatomic Ions – know red outlined ones for Exam 1 & rest for Exam 2.

Memorize: names, formulas & charges

IV. Nomenclature A. General Rule

Nomenclature: Name more + element & name more – element; change the ending to “ide” (for binary ionic compounds)

Examples of the “ide” ending:

Atom Anion Name

Chlorine Cl1- Chloride

Oxygen O2- Oxide

Fluorine F1- Fluoride

Sulfur S2- Sulfide

Nitrogen N3- Nitride

Iodine I1- Iodide

Bromine Br1- Bromide

Phosphorus P3- Phosphide

IV. Nomenclature B. Ionic Compounds

Name + element, then - element & change the ending to “ide.”

Exceptions: a) If have a polyatomic ion (pai), then name it; in the formula, place pai in ( ) if have MORE than one.

b) If have a variable charged metal, then give its charge with a Roman Numeral in parenthesis.

Examples:

NaCl = sodium chloride FeF3 = iron (III) fluoride

Li2O = lithium oxide LiNO2 = lithium nitrite

BaI2 = barium iodide Cu(CN)2 = copper (II) cyanide

AlF3 = aluminum fluoride Ca(NO3)2 = calcium nitrate

PbSO4 = lead(II)sulfate NH4I = ammonium iodide

IV. Nomenclature B. examples of names & formulas

NaCl = Sodium Chloride

BaI2 = Barium Iodide

Al2O3 = Aluminum Oxide

Ca(NO2)2= Calcium Nitrite

Cu2S = Copper(I) Sulfide

Sodium Hydroxide = NaOH

Ammonium Fluoride = NH4F

Iron(III)Sulfide = Fe2S3

Calcium Phosphate = Ca3(PO4)2

Beryllium Oxide = BeO

IV. Nomenclature C. Molecular Compounds, exception c.

If all atoms in a formula are nonmetals, then they combine with each other through sharing electrons (covalent bonds). They form molecules - no ions.

Name + element, then - element & change the ending to “ide.” Exception c = use prefixes to tell how many of each element is present (mono is optional).

When both elements are nonmetals (molecular compounds), then Name the +; the -; change ending to ide; & use prefixes of di, tri, tetra, penta, hexa, hepta, octa, nona, deca ( 2, 3, 4, 5, 6, 7, 8, 9 & 10 ).

Examples of non-metal compounds

CO = Carbon Monoxide

CO2 = Carbon Dioxide

NF3 = Nitrogen Trifluoride

N2F4 = Dinitrogen Tetrafluoride

PI5 = Phosporus Pentaoxide

Notes:

(1) organic compounds named with separate rules.

(2) diatomic molecules use the element name. Example: O2 = Oxygen

(3) Mono is optional.

IV. Nomenclature D. Examples continued

Need to go both ways: formula to name & name to formula. Need to know charges of ions and names/charges/formulas of polyatomic ions.

Carbon Dioxide = CO2

Barium Oxide = BaO

Calcium Hydroxide = Ca(OH)2

Sulfur Trioxide = SO3

Hydrogen Cyanide = HCN

Iron(III) Fluoride = FeF3

Barium Nitrite = Ba(NO2)2

IV. Nomenclature D. Examples Continued

NaF

CS2

NI3

BaI2

K3PO4

Iron(II) Oxide

Sodium Sulfate

Sodium Fluoride

Carbon Disulfide

Nitrogen Triiodide

Barium Iodide

Potassium Phosphate

FeO

Na2SO4

IV. Nomenclature E. Acids

- Acids are an important class of compounds that have their own set of names. Know before exam 2.

H2SO4 = Sulfuric AcidHNO3 = Nitric AcidH3PO4 = Phosphoric AcidHC2H3O2 = Acetic AcidHClO4 = Perchloric AcidHClO3 = Chloric AcidHClO2 = Chlorous AcidHClO = Hypochlorous AcidHCl = Hydrochloric AcidHBr = Hydrobromic AcidHF = Hydrofluoric AcidHI = Hydroiodic Acid

V. Nomenclature Summary

Know: - Metals/Nonmetals; Ionic/Molecular Compds

- Charges on ions

- Names, charges, formulas of polyatomic ions

- Prefixes: di, tri, tetra, penta, hexa

- Rules for nomenclature:

Name + then - & change ending to ide.

Exceptions:

- name polyatomic ion

- give charge on single multicharged ion with Roman Numeral in parenthesis

- use di, tri…. for two nonmetals

- Also be able to give formula from Name

VI. Chemical Equations A. Introduction

- A chemical equation is the representation of the rxn in terms of chemical formulas.

- Example: 2 Mg (s) + 1 O2 (g) -----) 2 MgO (s)

Notes: - “1” may not be shown; assume “1” if no # given.

- Mg & O2 are reactants.

- MgO is a product.

- ------) or = “goes to give” or yields.

- 2, 1, 2 are balancing coefficients.

- May add ΔH at the end for heat lost or gained.

- The subscripts of (g) (s) (l) (aq) indicate the physical state of the participants.

VI. Chemical Equations A. Introduction Continued

Notes Continued:

- The driving force in a chemical reaction is to produce products which are more stable than the reactants.

- How tell if a chemical reaction has taken place?

1) Gas produced 2) New solid, ppt, produced (ppt = precipitate)3) Light produced4) Heat may be lost or gained (temperature change)

- Can predict products with a) experience, with b) tables of solubility & c) with eqns of rxns that produce gasses.

- Have to balance by inspection & factor to simplest ratio.

VI. Chemical Equations B. Balancing

- Mass is neither lost or gained in a chemical reaction; so, we need to balance the equation without changing the identity of the reactants or products - change only balancing coefficients.

- Balancing is done by inspection, but for difficult equations it is best to balance the element which occurs the fewest times - first.

- Examples:

___Al + ___Cl2 -----) ___AlCl3

___Ca + ___H2O -----) ___Ca(OH)2 + ___H2

___HCl + ___Al(OH)3 -----) ___AlCl3 + ___H2O

___Ba(OH)2 + ___H3PO4 -----) ___Ba3(PO4)2 + ___H2O

atoms, molecules and ions. i. atoms & molecules a. atom examples - matter around us is made up...

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