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Exam 3 – Unit 2
 Complete
individually.
 No
notes. No talking.
 No
sharing calculators.
 When
 Begin
finished submit up front.
reading
 Chapter 3 pages 67-76 and
 Chapter 1 pages 16-20
 Complete Part 1 of Unit Packet
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Exam 3 Analysis

Average

Measurement and Calculations #1:
+____/25
%

Phases and KMT #2
+____/15
% (up from %)

Matter and Changes #3
+____/10
% (up from %)
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Unit 3 – Atoms: The Building Blocks
of Matter

Learning Objectives and Practice Packet
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Periodic Table Terminology
1. Periods: Horizontal row of elements in the
periodic table.
2. Group (family): A vertical column of
elements in the periodic table. (Have the
same number of valence electrons)
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Atomic Theory Continues…

We have come so far in our understanding of atoms.
Centuries of researching and countless scientists devoting
their lives to create the understanding of the atom today
(textbook concepts).

However it is NOT over. The more we understand about
atoms (how they work, their make up, etc…) the greater our
ability to advance science and technology in all aspects of
our lives (i.e. medicine).
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Check for Understanding

What is the charge of a proton?

Where would you find a proton in an atom?

What is the charge of an electron?

Where would you find an electron in an atom?

What is the charge of a neutron?

Where would you find a neutron in an atom?

How big is an electron compared to a proton?

How big is a neutron compared to a proton?
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Atoms:
The Building Blocks of Matter

Basic Laws (1790s)

Law of Conservation of Mass


Law of Definite Proportions


Mass is neither created nor destroyed during chemical reactions or
physical changes.
A chemical compound contains the same elements in exactly the same
proportions by mass regardless of the size of the sample or source of
the compound.
Law of Multiple Proportions

If two or more different compounds are composed of the same two
elements, then the ratio of the masses of the second element combine
with a certain mass of the first element is ALWAYS a ratio of small
WHOLE numbers.
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Atomic Theory - Video


John Dalton (1808)

English Schoolteacher

Proposed an explanation of the 3 basic laws (as mentioned
previously)
Dalton’s Atomic Theory (5 Statements)

All matter is composed of extremely small particles called atoms.

Atoms of a given element are identical in size, mass, and other
properties; atoms of different elements differ in size, mass, and
other properties.

Atoms cannot be subdivided, created, or destroyed.

Atoms of different elements combine in simple whole number
ratios to form compounds.

In chemical reactions, atoms are combined, separated, or
rearranged.
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
Modern Atomic Theory
Dalton’s Theory has been modified to fit new findings.
(Science is NOT static)


Example: Today we know that atoms are divisible into even smaller
particles

1. Subatomic Particles (protons, neutrons, and electrons)

2. Protons and Neutrons are made up of quarks
Important concepts

All matter is composed of atoms

Atoms of any one element differ in properties from atoms of another
element remain unchanged.
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The Atom

The smallest particle of an element that retains the chemical
properties of that element.

Consists of two regions:


Nucleus

1. Small region located at the center of an atom

2. Made up of at least one positively charged particle (proton)

3. Made up of usually one or more neutral particles (neutrons)
Region surrounding the nucleus – Electron Cloud

1. Very large compared to size of nucleus

2. Contains the negatively charged particles (electrons)
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Electron Discovery

JJ Thomson (1897)

Cathode-ray tubes

1. Electric current was passed through various gases at low
pressures.

2. The glow was caused by a stream of particles – cathode ray

3. The ray was deflected by a magnetic field and away from a
negatively charged object.

Concluded that all cathode rays are composed of identical
negatively charged particles (electrons)

Plum pudding model

1. Negative electrons spread throughout the positive charge of
the rest of the atoms.

2. Plums= electrons ; pudding=positive charge
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Electron Discovery Cont…

Robert Millikan (1909)

Measured the charge of the electron. (Charge to mass ratio)

Mass of electron = 9.109 x 10-31kg


Or about 1/2000 the mass of a hydrogen atom
Inferences from findings…

Because atoms are electrically neutral, they must contain a
positive charge to balance the negative electrons.

Because electrons have so much less mass than atoms, atoms must
contain other particles that account for most of their mass.
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Nuclear Atom

Ernest Rutherford (1911)

Bombarded a thin piece of gold foil with fast-moving alpha
particles (positively charged particles)

Concluded that deflected alpha particles must have experienced
some powerful force within the atom


The source of this force must occupy a very small amount of
space

Force must be caused by a very densely packed bundle of
matter with a positive electric charge (Nucleus)
Volume of nucleus was very small compared with the total volume
of an atom.

If the nucleus were the size of a pea, then the size of the atom
would be about the size of a football field.
+ Refresh: What are Atoms?
Atoms
are tiny particles that determine
properties of all matter.
Atoms
are the building blocks for
molecules.
Atoms
form elements.
Element: A
substance that cannot be
broken down into simpler substances by
chemical means.
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Parts of an Atom
Proton: A
subatomic particle that has
a positive charge and is found in the
nucleus of the atom.
Neutron: A
subatomic particle that
has NO charge and is found in the
nucleus of the atom.
Electron: A
subatomic particle that
has a negative charge and moves
around the outside of the nucleus.
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Label the Atom
Subatomic Particles
Particle Charge
Mass (kg)
Location
Proton
+1
1.67 x 10-27
nucleus
Neutron
0
1.67 x 10-27
nucleus
Electron
-1
9.11 x 10-31
Outside
nucleus
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Electron Orbital
-Electrons orbit the nucleus in orbital clouds.
-Electrons with different amounts of energy exist in
different energy levels.
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The Electron Cloud Model
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Electrons in each energy level
• Each energy level can hold a
limited number of electrons.
• The lowest energy level is the
smallest and the closest to the
nucleus.
Atomic Number
• The atomic
number of an
element is the
number of protons
in the nucleus of an
atom of that
element.
Formative Assessment

Read pages 77-87 and complete Part 2 of the unit packet!

Bring Questions!!
Formative Assessment Review
and Questions
Atomic Number
• The atomic
number of an
element is the
number of protons
in the nucleus of an
atom of that
element.
+ Charge of Atoms
 Atoms
are not charged even though they have
particles that contain charges.
 Atoms
are neutral because they have EQUAL
numbers of protons and electrons.
Ex: Helium Atom
Charge of 2 protons:
Charge of 2 neutrons:
Charge of 2 electrons:
Total charge of He atom:
+2
0
-2
0
Mass Number
 The
sum of the protons and neutrons in the
nucleus is the mass number of that
particular atom. (mass # = p + n)
+ Atomic Number vs. Mass Number
 Atomic
Number: Equal to the number of protons in the nucleus
of the atom.
(number of electrons = the number of protons)
 Mass
Number: Equal to the number of protons AND neutrons in
an atom’s nucleus.
 Average
Atomic Weight (below symbol on PT):
Weighted average of the atomic masses of the naturally occurring
isotopes of an element
 We will round this number to the nearest hundredth (TWO
decimal places)
 Example Oxygen’s average atomic mass is 15.9994 = 16.00

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Application
Iodine (I)
Iron (Fe)
Atomic number
____
Atomic Number
____
Atomic Mass
____
Atomic Mass
____
Number of Protons ____
Number of Protons ____
Number of Neutrons ____
Number of Neutrons____
Number of electrons ____
Number of Electrons ____
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Application
Nickel (Ni)
Radon (Rn)
Atomic number
____
Atomic Number
____
Atomic Mass
____
Atomic Mass
____
Number of Protons ____
Number of Protons ____
Number of Neutrons ____
Number of Neutrons____
Number of electrons ____
Number of Electrons ____
Isotopes
• Isotopes of an element have different mass
numbers because they have different
numbers of neutrons, but they all have the
same atomic number.
Isotope Examples
 Carbon
– 12
 6 protons
 6 electrons
 6 neutrons
 Carbon
– 13
 6 protons
 6 electrons
 7 neutrons
 Carbon
14
 6 protons
 6 electrons
 8 neutrons
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Example
6,7
Symbol
Li
3
Element
Name
Mass
Numbers
Lithium
Atomic
Number
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One More Example
107
Ag
47
Silver
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
Isotope Tables
Breaking it down…

To find the symbol – determine the atomic number of the element. This is the
number of protons

To find the protons- determine the atomic number of the element.

To find the electrons – equal to the number of protons of a neutral atom

To find neutrons: Mass Number – Atomic Number = Number of Neutrons

To find Mass Number: Atomic Number + Number of Neutrons = Mass
Number
+ Isotopes
+
Practice
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For Tomorrow…
 Complete
Isotope Worksheet and Part 3
worksheet!