Download atomic - Sammons Sci

Introduction to Chemistry
Atoms & Elements
Atom – are the fundamental particles of matter.

Means “uncuttable”, are about 1/10 of a billionth of a meter across.

Are composed of protons (p+), neutrons (nº) and electrons (e-)


Neutrons and protons are located in the center (nucleus)
Atoms are mostly empty space
- if an electron were basketball sized, the nucleus would be a car 20 miles
away.
Proton – means “first”

Relatively large, have a positive charge.

The # of protons in the nucleus determines what element the atom is (a.k.a. the atomic
number)
Neutron – means neutral



Approx. same mass as proton, no charge.
Determines whether an element is radioactive
Determines how much an element weighs
-

neutron + proton = atomic
mass
Forms of an element that have different numbers of neutrons are isotopes.
Electron


Small ( ~ 1/2000 mass of proton), negative charge.
Are attracted to protons

Move around in patterns known as orbitals.
- The orbital (or shell) is the distance from the nucleus.
-
The oribitals have names and are limited to the number of electrons they can hold.

The electrons in the outermost orbital are known as
valance electrons.
Nucleus



Contains protons and neutrons
Makes up most of the mass of an atom.
Small part of atoms volume.
Electron Cloud




Space in which electrons orbit the nucleus at high speed.
Large part of atoms volume.
Subdivided into Energy Levels or Shells.
Electrons occupy the lowest energy levels.
Modern Atomic Theory
Atoms are composed of 3 particles
Particle
Electron (e-, 0-1e)
Relative
Charge
-1
Proton (p+, 11H)
+1
Neutron (n0, 10n)
0
Actual Mass (kg)
Relative Mass
(amu) *
9.109 x 10-31
1.673 x 10-27
1.675 x 10-27
0.0005486
1.007276
1.008665
Mass
#
0
1
1
* Relative to 1/12 the mass of a carbon-12 atom
Protons (+) - Normally like charges repel each other, however when protons are extremely close
to one another, there is a strong
attraction between them
Nuclear Forces - short-range proton-proton, neutron-neutron, &
proton-neutron forces that hold the nucleus together.
Electrons (-) Move around nucleus in space, their - charge attracted to +charge of nucleus

e- do not orbit nucleus in well defined paths, but their location
cannot be

predicted at any given time
e- are said to be located in clouds (areas of probability) around the nucleus.
Atomic Models
Size of Atoms

Radius of e- cloud ranges from 40-270 picometer (10-12 m or 10-10 cm)

Radius of nucleus is  0.001 pm (very dense)
Atomic Number



All atoms are mode of p+, n0, & e-, but all are not the same element

Atomic number - the number of protons in the nucleus of each atom of a particular
element


On the periodic table, elements are placed in order by their atomic
Atomic number identifies an element

In an neutral atom, the number of p+ = number of e-
Atoms of different elements have different numbers of p+
Atoms of the same element have the same # of p+
#.
Mass Number - The total number of p+ & n0 in the nucleus of an atom.
Isotopes - atoms of the same elements with different masses (contains different numbers of
neutrons)



Isotopes occur naturally or are man-made, they do
not differ significantly in their
chemical behavior
Most elements are mixtures of isotopes, Tin has 10 - more than any other element
H is the simplest atom 1p+, but like many naturally occurring elements, H atoms can
contain different numbers of neutrons.
Isotope of H
p+
e-
n0
Abundance
Protium
Deuterium
Tritium
1
1
0
Most common (99.985% of H atoms)
1
1
1
0.015%
1
1
2
Radioactive, rare
To identify an isotope, you must know the name or atomic
the mass of the isotope.
Isotope of H
Protium
Deuterium
Tritium
p
n
+
0
1
1
1
0
1
2
number of element and
Mass
#
1
2
3
Nuclide - general term for any isotope of any element
There are 2 methods of designating isotopes
* H isotopes are unusual b/c they have distinct names
1. Mass # us written with a hyphen after the name of the element (uranium-
235)Hyphen Notation
2. Composition of nucleus is the isotope’s nuclear
Mass #  235
U
Atomic #  92
symbol
Number of n0 = mass # - atomic #
= 235 - 92
= 143
Sample Problems
1. How many p+, n0, & e- are there in an atom of chlorine-37?
p+ =
n0 =
(mass # - atomic #)
+
e=
(p = e in a neutral atom)
2. How many p+, n0, & e- are there in an atom of bromine-80?
p+ =
n0 =
e- =
3. Write the nuclear symbol for carbon 13:
4. Write the hyphen notation for element with 15 e- and 15 n0
5. Write the nuclear symbol for oxygen-16
6. Write the hyphen notation for element with 7 e- & 9 n0.
Relative Atomic Masses
Masses of atoms are very small
 An atom of oxygen-16 has a mass of 2.657 x 10-23g
For most chemical calculations relative atomic masses are used

The masses of all atoms are expressed in relation to the carbon-12 nuclide




Carbon-12 has been assigned a mass of exactly 12 atomic mass units (12 amu)
1 amu = 1/12 the mass of a carbon-12 atom
Atomic mass of any nuclide is determined by comparing it with the mass of carbon-12
Ex: Oxygen-16 has about 16/12 the mass of carbon-12 (15.994915 amu)
The mass # and relative atomic mass are quite close to each other, but not identical because
p+ and n0 masses are slightly >
1 amu and atomic masses include the mass of the e-.
Average Atomic Mass - the weighted (according to %) average of the atomic masses of
the naturally occurring isotopes of an element (See table 3-4)
Calculating Average Atomic Mass
Copper consists of 69.17% copper-63 (atomic mass of 62.92959 amu) & 30.83%
copper-65 (atomic mass of 64.927793 amu)
Convert % to decimal form and multiply the atomic mass
add the results

of each isotope &
0.6917(62.92959 amu) + 0.3083(64.927793amu) = 63.55 amu
Relating Mass to the Number of Atoms
Relative atomic mass makes it possible to know how many atoms of an element are present
in a sample of the element with a
measurable mass.
The Mole
The SI unit for amount
of substance
A mole (mol) is the amount of a substance that contains as many particles as there are
atoms in exactly 12


g of carbon-12.
The mole is a counting unit, just like a dozen is
The amount of substance that contains average number of particles
Avogadro’s Number
The number of particles in a mole is 6.0221367 x 1023
This means exactly 12 g of carbon-12 contains 6.0221367 x 1023 carbon-12 atoms
6.022 x 1023 (rounded) is the number of particles in exactly 1 mole of a pure
substance
Molar Mass
Question: Can we figure the approximate mass of 1 mol of He atoms?
Molar mass - the mass of 1 mol of a pure substance (written in g/mol)
The molar mass of an element is numerically equal to the element’s atomic mass
Ex: Molar mass of Li is 6.94 g/mol, Hg is 200.59 g/mol (rounded to 2 decimal places)
A molar mass of an element contains 1
mol of atoms
Answer: 4.00 g of He contains 1 mol of atoms
Gram/Mole Conversions
Chemists use molar mass as a conversion
factor
Ex: molar mass of He is 4.00 g/mol
How many g of He are in 2 mol of He
2.00 mol He (4.00 g He) = 8.00 g He
1 mol He
Mass of element in grams (molar
mass) x amount of element in mol (1/6.022 x
1023) = # of atoms in element
What is the mass in g of 3.50 mol of Cu?
3.50 mol C (63.55 g Cu) = 222 g Cu
1 mol Cu
Practice Problems
1. What is the mass in grams of 2.25 mol of the element iron?
2. What is the mass in grams of 0.375 mol of the element K?
3. What is the mass in grams of 0.0135 mol of the element Na?
4. What is the mass in grams of 16.3 mol of the element Ni?
Conversions with Avogadro’s Number
Use Avogadro’s Number to find number of atoms from amount in mol or amount in
mol from number of atoms.
How many moles of Ag are in 3.01 x 1023 atoms of Ag?
3.01 x 1023 atoms (
1 mol Ag
) = 0.500 mol Ag
23
6.022 x 10 Ag atoms
Practice Problems:
1. What is the mass in grams of 7.5 x 1015 atoms of nickel?
2. How many atoms of sulfur are in 4.00 g of sulfur?
3. What mass of gold contains the same number of atoms as 9.0 g of aluminum?
Periodic Table
Mendeleev arranged the elements based upon atomic mass and valance numbers to create
the periodic table.


Columns of elements are known as families or groups
- have similar but not identical properties
-
have the same valance
number
Rows are known as periods.
-
range from an active solid to an inactive gas.
Properties of Metals (located to the left of the zigzag line)
 Are good conductors of heat and electricity
 Shiny
 Have high melting points


ductile (can be made into wire)
Are malleable ( can be hammered)
Are
Properties of Nonmetals (located to the right of the zigzag line)
 Are poor conductors
 Have dull surfaces
 Have low melting points
 Are brittle and break easily
Metalloids – are located immediately on either side of the zigzag line.

-May have properties of metals and nonmetals
From left to right on the periodic table …
 elements on the left side lose e-, on the right gain e-.
 the amount of energy required to remove an e- increases


the atomic size decreases
the elements become less
The Periodic Law
metallic
Dmitri Mendeleev (Russian) – was the first to place atoms in order according to their
properties. (periodic table)
Credited with discovery of periodic law
Placed most elements in order of atomic mass and noticed repeating (periodic) trends
in this arrangement.
Did not exactly work arranging according to atomic mass
1. Why could most elements be arranged according to mass, but a few could not?
2. What was the reason for chemical periodicity?
Henry Moseley (English) – discovered 40 years later that elements in periodic table fit
into patterns better when arranged to nuclear charge (number of protons  atomic
number)
 atomic number is basis for periodicity, not
mass
Periodic Law – the physical and chemical properties of the elements are periodic
functions of their atomic numbers.
When arranged by increasing atomic numbers, elements with similar properties appear at regular
intervals.
Periodic Table – an arrangement of the elements in order of their atomic numbers so that
elements with similar properties fall in the same column, or group.
Electron Configuration and the Periodic Table
Electron Configuration & Periodic Table
The valence electrons for elements in a group (column) will have a similar electron
configuration for the valence electrons. The only difference will be which shell holds the
valence electrons.


In general, the group number equals the number of valence electrons.
The electron configuration can be determined by the periodic table because the shell and

subshell information is embedded in the table.
Blocks of the periodic table, corresponding to filling the different kinds of orbitals.
Beginning at the top left and going across successive rows of the periodic table
provides a method for remembering the order of orbital filling: 1s --> 2s --> 2p --> 3s -->
3p --> 4s --> 3d --> 4p, and so on.
S-Block Elements: Groups 1 & 2
Group 1 – Alkali Metals (Li,
Na, K, Rb, Cs, Fr)
Group electron configuration: ns1
If pure, they are silvery & soft enough to cut with a knife
Very reactive, lose electrons from outer shell to form stable
structure
not found as free elements in nature
React strongly with nonmetals
React strongly with H2O to form H2(g) and alkali solutions
Stored in kerosene so they do not react with air or moisture
From top  bottom of column, alkali metals have successfully lower melting points.
Group 2 – Alkaline Earth Metals (Be, Mg, Ca, Sr, Ba, Ra)
Group electron configuration: ns2
Have a pair of electrons in outer shell
Harder, more dense,
stronger than alkali metals, & have higher melting points
Still too reactive to be found as free elements in nature
Hydrogen & Helium
H has 1 electron in outer energy level, but does
not have same properties as Group 1
He has 2 electrons in outer energy level, but is in Group 18 (Noble Gases) and has a
filled outer energy level
Without looking at Periodic Table:
1. What is the group, period & block of elements with electron configuration [Kr]5s2?
2. What is the group configuration for Group 1 elements?
3. Write the electron configuration for Group 1 element in 5th period.
Using the Periodic Table:
1. What is the above element’s noble gas notation?
2. How does reactivity of this element compare to that of element in group 2 of the same period?
d-Block Elements: Groups 3-12 are Transition Metals
Have typical metallic properties
Good conductors of heat & electricity
Have a high luster (shiny)
Less reactive than Groups 1 & 2
Palladium, platinum & gold are among least reactive in nature
Do not easily form compounds, exist as free elements
p-Block Elements: Groups 13-18 Main Group
Properties vary greatly
Include nonmetals, metals & metalloids

B, Si, Ge, As, Sb, Te
Group 17 Halogens
Are the most reactive nonmetals
React vigorously with most metal to form salts
Have 7 electrons in outer shell
Metalloids are along zigzag line, brittle metals with properties of both metals and
nonmetals.
Metals of p-block are generally harder & more
dense than d-block metals
All except bismuth are reactive enough to be compounds in nature, but once they are
obtained as free metals, they are stable in air
F-Block Elements: Lanthanides & Actinides
Lanthanides – 14 elements between lanthanum & hafnium, shiny metals similar in
reactivity to Group 2 metals
Actinides – 14 between Actinium & Rutherfordium, all radioactive
First four (Thorium  Neptunium) are found in nature, the rest are man-made
Periodic Table Trends
Electron Configuration & Periodic Properties
Atomic Radii
Atomic Radius – one-
half of the distance
between the nuclei of
identical atoms that
are bonded together
Atomic size
decreases as you
move to the right across a period and increases as you move down a column
in the periodic table.

Atomic size increases down a column because orbitals of higher quantum number (n)
have their maximum probability farther from the nucleus.

Atomic size decreases from left to right in the periodic table because the greater number
of protons in the nucleus will exert a greater attraction on the electrons, pulling

them closer to the nucleus.
Electrons moving across the nucleus do not experience the same nuclear attraction; those
electrons closer to the nucleus experience a greater force than those that are farther
away.

The nuclear charge actually "felt" by an electron is called the effective nuclear
charge, Zeff.


Ideally, the size of an atom defined by the edge of its orbital  varies according to
conditions in which atom exist
Period Trends
Radii decreases from left to right across a period due to increasing positive charge
on nucleus

Group Trends
Radii increases from top to bottom of group due to addition of e- in higher main energy
level
1. Describe how the atomic radius changes
within a period.
Gets smaller b/c pulls tighter
2. Describe how the atomic radius changes
between periods.
Gets larger, more orbitals
3. Why is radius of Ga less than Al?
More opposite charges & d block
Ionization Energy
An e- can be removed from any atom if enough energy is supplied
A + energy  A+ + e-

A+ represents an ion - an atom with a positive or negative charge

ionization - any process that results in the formation of an ion
Period Trends (general)


ionization energy increases
from left to right across a period
this is due to increasing nuclear
charge, greater charge =
stronger attraction of e
Group Trends (general)
 ionization energy
decreases moving down a

group
this is because the e- are
further away from the attraction of the positively charged nucleus
Electron Affinity

Electron affinity - the energy change that occurs when an e- is acquired by a neutral
atom

most atoms release energy when they acquire an eA + e-  A- + energy
Partial periodic table - metals tend to lose electrons to attain noble gas electronic
configuration. Non-metals tend to gain electrons to attain noble gas electronic
configuration

The general trend is for the electron affinity to become increasingly negative
(stronger
halogens.
binding of an electron) as we move across each period toward the
some atoms must be “forced” to gain an e-