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Matter and Properties

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					                                                      Unit 1
                                        The Modern Period Table of Elements

An element is a pure substance that cannot be broken down into simpler substances by chemical
means.
    composed entirely of only one kind of atom

   Elements that conduct heat/electricity, are malleable, ductile, and lustrous are classified as metals.

   Elements that are non-conductive and brittle as solids are classified as nonmetals.

   Elements that have share metallic and nonmetallic properties are classified as metalloids.

Modern periodic law states that when elements are
arranged in order of their increasing atomic number,
their properties show a periodic recurrence and gradual
change.

   A group is a vertical column in the periodic table
    with elements that share similar chemical
    properties.

   A period is a horizontal row of elements whose
    properties change from metallic on the left to
    nonmetallic on the right.
                alkaline earth metals




                                                                                    noble gases
                      alkali metals




                                                                                    halogens




                                          transition metals




                                                       lanthanides
                                                       actinides
                                 Groups and Series of Elements

Representative elements (most closely follow periodic law)
alkali metals           soft, silver coloured metals                       Li
(Group 1)               solids at room temperature                         Na
                        exhibit metallic properties                        K
                        react violently with water
                        react with halogens
                        stored under oil to prevent reaction with air
alkaline earth          light, reactive metals                             Mg
metals (Group 2)        solids at room temperature                         Ca
                        exhibit metallic properties                        Sr
                        react violently with acids
                        form oxide coatings when exposed to air
halogens                solids, liquids, or gases                          F
(Group 17)              exhibit nonmetallic properties                     Cl
                        extremely reactive with hydrogen and metals        Br
noble gases             gases at room temperature                          He
(Group 18)              extremely unreactive                               Ne
                        Kr, Xe, and Rn reluctantly form compounds with F   Ar
                        radon is radioactive
transition metals       strong, hard metals with high mps                  Fe
                        conduct heat and electricity                       Cu
                        variable reactivity                                Ag
                        many react with oxygen to form oxides              Zn
                        some react with acids to form hydrogen gas         Co

lanthanides             elements with atomic nos. 57-70                    Ce
actinides               elements with atomic nos. 89-102                   U
                                   Developing a Model of the Atom

Empirical knowledge comes directly from observations and experimentation.

Theoretical knowledge is knowledge based on ideas created to explain observations.

   A theory is a comprehensive set of ideas that attempts to explain a law or related observations.
   A model is a mental or physical representation of a theoretical concept.

The smallest particle of an element that has all the properties of that element is an atom.
The small, positively charged centre of an atom is the nucleus.

   A proton is a positively charged subatomic particle in the
    nucleus of an atom.
   An electron is a negatively charged subatomic particle.
   A neutron is an uncharged subatomic particle in the nucleus
    of an atom.


                 Understanding Atomic Mass

Atomic number (Z) is the number of protons present in the nucleus of an atom of a given element.
 The number of electrons equals the number of protons in an atom.

Atomic mass (A) is the sum of the number of protons and neutrons present in the nucleus of an atom.

                         number of neutrons = mass number – atomic number

                                             N=A–Z
N represents the number of neutrons. A collective term for protons and neutrons, the elementary
particles found in the nucleus, is nucleons.

Isotopes are atoms of an element that differ from each other only in the number of neutrons in their
nuclei.

                                   mass number              A                 14              12
                                 atomic number              ZX                  C              C
                                                                               6              6


                                A naturally or artificially produced element that is capable of
                                 spontaneously emitting radiation in the form of particles and/or
                                 gamma rays is a radioisotope.
                                               Nuclear Chemistry

The majority of nuclei are unstable and exhibits radioactivity.
 radioactive means capable of spontaneously emitting radiation in the form of particles and/or
   gamma rays
    radioactive decay

Each type of unstable nuclei has a characteristic rate of decay, which varies from fractions of a second
to billions of years.

Radioactive Decay
When nuclide of one element decays, it changes into nuclide of a different element.
 transmutation

The decaying nuclide is called the parent, and the product nuclide is called the daughter. Nuclear
decay produces a nuclide of lower energy as the excess energy is emitted as radiation.

During any transmutation event, the total Z (number of protons) and total A (sum of protons and
neutrons) of reactants must equal those of the products.

                                     TotalA
                                     Total Z   Reactants    TotalA
                                                             Total Z   Products


                             Modes of Radiation and Radioactive Decay

   Alpha particles (symbolized  or 4 He ) are dense, positively charged particles identical to helium
                                     2
    nuclei.
     released during alpha decay
                         92 U   90Th  2 He
                        238     234     4


                  reactant  product +  particle
                    parent  daughter + (2n+2p)

   Beta particles (symbolized  or 1 e ) are negatively charged
                                         0


    particles identified as high-speed electrons.
     released during beta decay
                              6C     7 N   1e
                             14      14       0


                     reactant  daughter +  particle
                               0 n  1p   1e
                               1     1      0




   Gamma rays (symbolized  or 0  ) are very high-energy photons, about 105 times as energetic as
                                0
    visible light.
     released during gamma emission (accompanies most other types of decay)

                                        U  234 Th  4 He  2 0 
                                       238
                                        92    90     2        0

                                  parent  daughter +  particle + 2 gamma rays
                                         Measuring Radioactive Decay

All radioactive material decays at a characteristic rate, regardless of the chemical substance in which
they occur.
 the rate of decay is called the activity and can be measured with a Geiger counter
 the SI unit of activiy is the becquerel (Bq), defined as one disintegration per second (1 Bq = 1 d/s)
     the curie (Ci) is equal to the number of nuclei disintegrating each second in 1 g of radium-226


1 Ci = 3.70 x 1010 Bq

Half-life (t1/2) is the time it takes for half of the nuclei present
to decay.
 The amount of nuclei remaining is halved after each half-
    life.
 Every radioisotope has its own characteristic value for
    half-life, and can be measured in seconds (s), minutes
    (min), hours (h), or years (a).

        14
         6   C       14
                       7   N    0
                                1   e     t 1/ 2  5730 a


                                             Radioisotopic Dating

   application of radioisotopes naturally present to determine the age of an object
   the accuracy of the method diminishes after 6 half-lives and specific radioisotopes can be used for
    different substances and various lengths of time

A technique that uses radioactive carbon-14 to identify the date of death of once living material is
called carbon-14 dating.
 14C : 12C ratio is determined based on radioactivity of the sample, from which the elapsed time can
    be calculated

                t                              A is the measured activity
A  A0  2     t1/2                             A0 is the normal activity
                                                t is the time elapsed
                                                t1/2 is the half-life

   potassium-40 and uranium-238 can be used to date non-living materials with much greater ages
    (e.g. rocks)
                                            Atomic Spectra

Continuous spectrum is the pattern of colours observed when a narrow beam of white light is passed
through a prism or spectroscope.
 contains all wavelengths of visible light

Line spectrum is the pattern of distinct lines, each of which corresponds to light of a single
wavelength.
 produced when light consisting of only a few distinct wavelengths is passed through a prism or
   spectroscope

When the atoms of elements are energized, they emit characteristic patterns to their line spectra.
 each line corresponds to a specific quantity of energy as an electron returns to its original energy


                                       Modern Atomic Theory

The orbitals of an electron are actually regions of fixed energy in which an electron is allowed to move
and orbit the nucleus.
 energy levels (or energy shells)


Features of the First Four Shells
          Number of         Maximum number of electrons per        Maximum Number of electrons per
 n
          subshells                    shell                                    subshell
1    1                     2                                      (2)
2    2                     8                                      (2 + 6)
3    3                     18                                     (2 + 6 + 10)
4    4                     32                                     (2 + 6 + 10 +14)


                          Ground state is the lowest energy level that an electron can occupy.

                          Transition is a movement from one energy level to another.

                          Valence electrons are those electrons in the highest energy level of an atom.
                                     Trends in the Periodic Table

Atomic Radius
 the distance between the nuclei of bonded or adjacent atoms of
   the same element divided by two

   down a group, the atomic radius gets larger (adds on extra
    shell)
   across a period, the atomic radius gets smaller (more electrons
    in valence shell, attracts to nucleus)



Ionic Radius
                                          a measurement of the size of an ion
                                            positive ions are smaller than atoms (same positive
                                              charge in nucleus pulls electrons tighter)

                                              negative ions are larger than atoms (greater number of
                                               electrons, more negative charge, nucleus cannot hold as
                                               tightly)




                   Trends in the Periodic Table

Ionization Energy
 the amount of energy required to REMOVE an electron from
   an atom or ion in the gaseous state

                        X(g)           +        IE1          X+(g)   +      1e-

Moving down a group,
 increase in number of orbitals
 increase in atomic radius
   since the electrons are further from the nucleus, they are more loosely held and easier to
      remove

The first ionization energy decreases down a group.

Moving across a period,
 increase in number of protons in the nucleus
 decrease in atomic radius
   since the electrons are closer to the nucleus, they are more tightly held and harder to remove

The first ionization energy increases across a period.
Second Ionization Energy

By removing a second electron from the cation formed in the gaseous state, the second ionization
energy is measured.



             + IE1                      +     IE2             +       IE3
    atom                     +1 ion                      +2 ion
   after the removal of first electron, ionic radius is smaller as electrons are held more tightly
   more energy is required to remove second electron
     second ionization energy is larger


                                 IE1 < IE2 << IE3 <<< IE4


Electron Affinity
 the energy released when an electron is accepted by an atom in the gaseous state

                X(g)            +       1e-                            X-(g)   +      EA1

The stronger an atom holds its valence electrons, the more likely that atom will want to gain another
electron.

[EXCEPTION: noble gases already have a full octet]

                                              As atomic radius increases, the atom has looser hold on
                                              electrons, the electron affinity decreases.
                                               down a group, EA decreases

                                              As atomic radius decreases, the atom has tighter hold on
                                              electrons, the electron affinity increases.
                                               across a period, EA increases
                                           Activity Series

Reactivity is a chemical property of an element that indicates the tendency of the element to form a
compound.
 an element that forms compounds easily has high reactivity
 an element that does not form compounds has low reactivity

An activity series is a list of elements in order of their reactivity, based on evidence gathered from
single displacement reactions.


          HIGHEST REACTIVITY
                 lithium
                potassium
                 barium
                 calcium
                                                                 HIGHEST REACTIVITY
                  sodium
                magnesium                                                fluroine
                aluminum
                                                                         chlorine
                    zinc
                    iron                                                 bromine
                  nickel                                                  iodine
                     tin
                    lead
                                                                 LOWEST REACTIVITY
                (hydrogen)
                  copper
                    silver
                     gold

          LOWEST REACTIVITY
                                            Chemical Bonds

Ionic compounds are pure substances
composed of ions and formed from a metal
and a nonmetal.
 Ionic bonds are a type of chemical bond
   created by the electrostatic attraction
   between positive and negative ions.

Octet rule states that representative elements
want a stable octet consisting of a full shell
of eight electrons in their outer energy level.
 metallic elements lose electrons (positive
    ions)
 nonmetallic elements gain electrons
    (negative ions)

Molecular compounds are pure substances composed of two
or more nonmetals.
 Covalent bonds are a type of chemical bond created by an
    attractive force between two atoms when they share
    electrons.

   nonmetallic elements want to gain electrons and are
    involved with covalent bonding

                                       Properties of Compounds

The simplest whole number ratio of ions in an ionic compound is called the formula unit.
 ions actually form a crystal lattice – a regular, ordered arrangement of ions

           Properties                                Ionic                     Molecular

          Melting point                              High                         Low

     Electical conductivity
                                             No conductivity                 No conductivity
         in solid state
                                                                             No conductivity
          in liquid state                         Conductivity


      Consistency of solid                        Hard, brittle           Soft, waxy or flexible

Intramolecular force is the attractive force between atoms and ions within a compound.
 strong intramolecular force in ionic solids

Intermolecular force is the attractive force between molecules.
 experienced by molecules of molecular compounds
                                               Lewis Symbols

   a representation of an atom or ion, made up of the chemical symbol and dots indicating the number
    of electrons in the valence energy level


                 H                                                                            He
                 Li       Be                   B C            N       O                  F    Ne
                 Na Mg                         Al Si P                S                  Cl Ar
In the formation of ionic compounds, show the movement of electrons with a full arrow:
                                  .             .                    +2    ..       -2
                                                                 [Mg] [ O ]




                                                                               ..
                                 Mg        +   O                        ..




                                                                          ..
                                                ..
                                      .



                                               ..
                                               .



In the formation of molecular compounds, show the movement of electrons with a half arrow:
                            ..            ..         ..     ..             ..            ..
                           Cl + Cl                   Cl . . Cl             Cl Cl
                         ..




                                                    ..




                                                                          ..


                                                                                         ..
                                          ..




                                                             ..
                            ..   ..                   ..     ..             .. ..
                                      .
                              .




   the covalent bond is shown with a solid line


         Drawing Lewis Structures and Structural Formulae for Molecular Compounds

Step 1         Place the element with the highest bonding capacity in the central position and arrange
               the symbols of the other elements around it.
Step 2         Add up the number of valence electrons available in each of the atoms. For polyatomic
               ions, add one electron for each unit of negative charge, or subtract an electron for each
               unit of positive charge.
Step 3         Place one pair of electrons between each adjacent pair of electrons, forming single
               covalent bonds.
Step 4         Place pairs of the remaining valence electrons as lone pairs on the peripheral atoms
               only.
Step 5         If octets are not complete, move lone pairs into bonding position between those atoms
               and the central atom until all octets are complete.
Step 6         If the peripheral atoms all have complete octets and there are pairs of electrons
               remaining, place these electrons as lone pairs on the central atom.
Step 7         To give the structural formula, remove the dots representing the lone pairs and replace
               bond dots with dashes.
Step 8         If you are representing a polyatomic ion, place brackets around the entire structure and
               write the charge outside the brackets.
                                     Coordinate Covalent Bond

Molecular compounds often have pairs of electrons that are not involved with bonding called lone
pairs.

   Lone pairs can be used to form covalent bonds, in which one atom donates both electrons to be
    shared.
   coordinate covalent bonds


                                                                 +

                      A            B               A       B
                       ..

                                                                     coordinate
                                                                     covalent
                           -
                      A            B               A       B         bond
                       ..




The ammonium ion (NH4+) is an example of coordinate covalent bond.

                                                                              +
                         H+            ..                        H
                               H       N     H
                                                           H     N      H
                                       H
                                                                 H


                                   Identifying Coordinate Bonds

   Draw the Lewis structure for the compound.
   Count the number of bonds between the central atom and the peripheral atoms.
   Compare the number of bonds to the bonding capacity of the central atom.
   If the number is more than the bonding capacity, one or more of the bonds is coordinate covalent.
   To identify which ones, remove the peripheral atoms one at a time.
   If you can do this and leave the central structure with a complete octet, you have identified
    coordinate covalent bonds.
                                             Polar Bonds

Within a covalent bond, the electronegativity (Pauling) of each bonding atom affects the type of bond.
 electronegativity is a property of an atom within a covalent bond describing its ability to attract
   the electrons in the bond to its nucleus

Nonpolar covalent bond results from a zero difference in electronegativity between the bonded atoms
and has equal sharing of bonding electrons.

Polar covalent bond results from a difference in electronegativity between the bonded atoms and the
electrons are not shared equally.
 results in partial positive and negative ends of the bond called dipoles (+ and -)

When the electronegativity difference is high, the bonding pair of electrons is mostly with one
atom/ion resulting in an ionic bond.
 the greater the difference in electronegativity, the more ionic character of the bond



                                           Polar Molecules

Both the polarity of the bonds and the shape of the molecule determine if the molecule is polar or
nonpolar.

   A nonpolar molecule has either all nonpolar bonds or polar bonds whose bond dipoles cancel,
    such that there is no positive/negative ends.
     Symmetrical molecules with polar bonds are generally nonpolar since the sum of the bond
       dipoles is zero.

   A polar molecule has polar bonds with dipoles that do not cancel and create positive and negative
    ends of the molecule.
                                          Intermolecular Forces

    the force of attraction and repulsion between molecules

        generally much weaker than covalent and ionic bonds (intramolecular forces)

1.   dipole-dipole forces

        a force of attraction between polar molecules where the oppositely charged dipoles attract one
         another
        the greater the polarity of the molecule, the greater the strength of the dipole-dipole force

2.   hydrogen bonds

        special dipole-dipole forces involving N-H, O-H, and F-H bonds
        highly polar bonds and molecules, resulting in strong intermolecular forces between the
         hydrogen atom and the lone pair of N, O, or F

3.   London dispersion forces

         weak intermolecular forces in nonpolar molecules
         due to the simultaneous attraction of the elecrons of one molecule by the positive nuclei in the
          surrounding molecules
          by increasing the number of electrons in the molecule, the strength of London forces
             increases

				
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