1. Unit 1 Section 4 Atomic Structure

2. Objectives Differentiate between three types of subatomic particles
Define isotope and calculate the atomic mass
Explain early models of the atom

3. Sections 1.4.1 Defining the Atom 1.4.2 Structure of the Nuclear Atom
1.4.3 Distinguishing Between Atoms

4. 1.4.1 and Defining the Atom
Vocabulary
Atom
Dalton’s atomic theory
Atom– is the smallest particle of an element that retains its identity in a chemical reaction.

5. 1.4.3 Distinguishing Among Atoms
Vocabulary
Atomic number
Mass number
Isotopes
Atomic Mass Unit
Atomic mass
Periodic table
Period
Group
Atomic number– the number of protons in the nucleus of an atom of an element
Mass number- the total number of protons and neutrons in the nucleus of an atom
Isotopes- atoms of the same element that have the same atomic number but different atomic masses due to a different number of neutrons
Atomic mass unit- a unit of mass equal to one-twelfth the mass of a carbon-12 atom
Atomic mass- the weighted average of the masses of the isotopes of an element
Periodic table- an arrangement of elements in which the elements are separated into groups based on a set of repeating properties
Period- a horizontal row of elements in the periodic table
Group- a vertical column of elements in the periodic table; the constituent elements of a group have similar chemical and physical properties

6. Experiencing Atoms
Atoms are incredibly small, yet they compose everything
Atoms are the pieces of elements
Properties of the atoms determine the properties of the elements
If every atom within a pebble were the size of the pebble itself, then the pebble would be larger than Mt. Everest (~29,000 ft)

7. Atomic Structure and the Periodic Table of Elements
Protons: positively charged; located in the nucleus Neutrons: neutral charge; located in the nucleus Electrons: negatively charged; located around the nucleus Nucleus: contains protons and neutrons; center of the atom; positively charged Atomic Mass: mass of element; sum of protons and neutrons Atomic number: # of protons  also equal to # of electrons

8. C 6 12.01 Symbols Carbon Find the Atomic number Mass Number
number of protons
number of neutrons
number of electrons
Element name 6
12.01 C
Element symbol
Carbon
Element name
Atomic
number

9. Unit 1 Section 4 Atomic Structure
Atomic Number
The atomic number identifies an element.
*Elements are different because they contain different numbers of protons.
Remember that Atoms are electrically neutral.
In an atom, protons = electrons

10. Mass Number Mass number = protons + neutrons How many neutrons??
The number of neutrons in an atom is the difference between the mass number and atomic number.

11. Atoms of the First Ten Elements
Interpret Data
For each element listed in the table below, the number of protons equals the number of electrons. (Protons = Electrons)
Atoms of the First Ten Elements
Name
Symbol
Atomic number
Protons
Neutrons
Mass number
Electrons
Hydrogen H 1
Helium He 2 4
Lithium Li 3 7
Beryllium Be 5 9
Boron B 6 11
Carbon C 12
Nitrogen N 14
Oxygen O 8 16
Fluorine F 10 19
Neon Ne 20
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12. Do Now
1. Fill in the chart.
2. The atomic number represents what two particles?
Element Name
Atomic Mass
Atomic #
Protons
Electrons
Neutrons
Fluorine (F)
Bromine (Br) 19 9 9 9 10 80 35 35 35 45

14. Isotopes
4.3
Isotopes are atoms that have the same number of protons but different numbers of neutrons.
Because isotopes of an element have different numbers of neutrons, Isotopes also have different mass numbers.

15. Isotopes
4.3
Despite these differences, isotopes are chemically alike because they have identical numbers of protons and electrons.
Neon-20, neon-21, and neon-22 are three isotopes of neon, a gaseous element used in lighted signs. Comparing and Contrasting How are these isotopes different? How are they similar?

16. Isotopes For example, hydrogen has 3 isotopes:
Atomic #
Atomic Mass
Proton
Neutron
Electron
Hydrogen-1
Hydrogen-2
Hydrogen-3
Note that the correct way to write an isotopes is to write the name, followed by the mass number.

17. Isotopes A – Z = n (number of neutrons) X = the symbol of the element
Z = the atomic number (# of protons)
A = the mass number (# of protons and neutrons)
A – Z = n (number of neutrons)

18. Isotopes – An Example F = the symbol for fluorine
9 = atomic number
19 =
Proton # =
Neutron # =
F = the symbol for fluorine
9 = atomic number
21 =
Proton # =
Neutron # =
the mass number
the mass number 9 9 10 11

19. Why are atoms with different numbers of neutrons still considered to be the same element?
-Despite differences in the number of neutrons, isotopes are chemically alike.
-They have identical numbers of protons and electrons, which determine chemical behavior.
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20. Objectives
To calculate average atomic mass of elements.

21. Atomic Mass Unit – How is it calculated?
The value shown in the periodic table is the average atomic mass
It is a weighted average of the atomic masses of naturally occurring isotopes

22. 4.3
Average Atomic Mass
To calculate the average atomic mass of an element, multiply the mass of each isotope by its relative abundance, expressed as a decimal, and then add the products.

23. Example 1 – Calculating Average Atomic Mass
Carbon has two stable isotopes: carbon-12 and carbon-13.
Isotope
Mass (amu)
Relative Abundance
Calculations
Carbon-12
amu
98.89 %
Carbon-13
amu
1.11 %

24. Average Atomic Mass of carbon = (12. 000 amu x 0. 9889) + (13
Average Atomic Mass of carbon = ( amu x ) + ( amu x )
= ( amu) + (0.144 amu)
= amu

25. Example 2 – Isotopes of Chlorine
Cl-35 has an amu of with an abundance of %
Cl-37 has an amu of with an abundance of %
Calculate the average atomic mass.
Isotope
Mass (amu)
Relative Abundance
Calculations
Cl-35
75.771%
Cl-37
24.229%

26. Example 3 – Try
Calculate the average atomic mass of Europium given the percent abundances and relative masses.
Relative masses
Percent abundance
Europium-151
151
48.03%
Europium-153
153
51.97%

28. DO NOW What is an isotope?
Give two examples of isotopes of the SAME ELEMENT.
3. How many isotopes of CARBON are there?

29. Objectives To learn how a formula describes a compound’s composition.
To study the atom’s structure

30. The Elements
Substances that cannot be broken down by simple chemical means
118 elements known

31. Element Abundance
Most abundant elements in the universe: hydrogen H 60% helium He 37 Most abundant elements in the entire earth: iron Fe 35% silicon Si 15% oxygen O 30% Most abundant element in earth’s crust: oxygen O 49% aluminum Al 7% silicon Si 26%

32. Element Abundance Naturally radioactive elements:
uranium U
radium Ra
radon Rn
Polonium Po
Elements most abundant in human body:
oxygen O nitrogen N
hydrogen H phosphorus P
carbon C calcium Ca

33. Diatomic Elements
Diatomic elements – elements that occur naturally paired as two atoms per molecule
- 7 diatomic elements
H2 Cl F2
N2 Br2
O2 I2
GEN-U-INE

34. Names and Symbols of the Most Common Elements

35. Scientists of the 18th century learned that:
Pure substances are either elements or combinations of elements called compounds.
Law of Constant Composition
A given compound always contains the same proportions (by mass) of the elements.

36. Law of Constant Composition
A given compound always has the same composition, regardless of where it comes from.
Water always contains 8 g of oxygen for every 1 g of hydrogen.

37. Chemical Formulas Describe Compounds
Compound – distinct substance that is composed of the atoms of two or more elements and always contains exactly the same relative masses of those elements.
Chemical Formulas – expresses the types of atoms and the number of each type in each unit (molecule) of a given compound.

38. Rules for Writing Formulas
Symbol  Tells which atoms are present in compounds
Subscript  The number of each type of atom
Coefficient  (number in front) tells the number of molecules
Coefficient
Copy example above

39. 4.2 Structure of the Nuclear Atom
Vocabulary
Electron
Cathode ray
Proton
Neutron
Nucleus
Cathode ray- a stream of electrons produced at the negative electrode (cathode) of a tube containing a gas at low pressure

40. Subatomic Particles What are three kinds of subatomic particles?
4.2
Subatomic Particles
What are three kinds of subatomic particles?
Three kinds of subatomic particles are electrons, protons, and neutrons.

41. Early Models of the Atom
John Dalton ( )
Dalton’s Atomic Theory
Stated all elements are composed of tiny indivisible particles called atoms  atoms cannot be broken down further.
Elements can form compounds

42. Protons Eugen Goldstein (1850–1930)
Observed a cathode-ray tube and found rays traveling in the direction opposite to that of the cathode rays.
He concluded that they were composed of positive particles.
Goldstein discovered protons

43. Electrons English physicist J.J. Thomson
discovered the electron (1897)
Experiment: Cathode ray tubes
Thomson performed experiments that involved passing electric current through gases at low pressure.

44. Thomson’s Plum Pudding Model
Reasoned that the atom might be thought of as a uniform “pudding” of positive charge with enough negative electrons scattered within to counterbalance that positive charge.
Believed the electrons are evenly scattered throughout a cloud of positive charge

45. Electrons
The U.S. physicist Robert A. Millikan (1868–1953) discovered the charge of an electron
Using this charge and Thomson’s charge-to-mass ratio of an electron, Millikan calculated an electron’s mass.
An electron has one unit of negative charge, and its mass is 1/1840 the mass of a hydrogen atom.

46. Neutrons Physicist James Chadwick (1891–1974) Discovered the neutron
Neutrons are subatomic particles with no charge but with a mass nearly equal to that of a proton.

47. Subatomic Particles
4.2
Table 4.1 summarizes the properties of electrons, protons, and neutrons.

48. The Atomic Nucleus
4.2
Born in New Zealand, Ernest Rutherford was awarded the Nobel Prize for Chemistry in His portrait appears on the New Zealand $100 bill.

49. Ernest Rutherford Ernest Rutherford discovered the nucleus (1911)
Stated protons were inside of the nucleus
Gold-Foil Experiment
Radioactive particles “shot” through gold foil
A majority of particles passed straight through foil
A small fraction of particles bounced off gold foil at large angles or bounced straight back

50. The Atomic Nucleus Rutherford’s Gold-Foil Experiment
4.2
Rutherford’s Gold-Foil Experiment
Rutherford’s gold-foil experiment yielded evidence of the atomic nucleus. a) Rutherford and his coworkers aimed a beam of alpha particles at a sheet of gold foil surrounded by a fluorescent screen. Most of the particles passed through the foil with no deflection at all. A few particles were greatly deflected. b) Rutherford concluded that most of the alpha particles pass through the gold foil because the atom is mostly empty space. The mass and positive charge are concentrated in a small region of the atom. Rutherford called this region the nucleus. Particles that approach the nucleus closely are greatly deflected.
Be able to explain Gold Foil Experiment

51. The Atomic Nucleus
4.2
In the nuclear atom, the protons and neutrons are located in the nucleus.
The electrons are distributed around the nucleus and occupy almost all the volume of the atom.

52. Rutherford’s model

53. How can there be different varieties of atoms?
Just as there are many types of dogs, atoms come in different varieties too.
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54. Key Concepts
Elements are different because they contain different numbers of protons.
Because isotopes of an element have different numbers of neutrons, they also have different mass numbers.
To calculate the atomic mass of an element, multiply the mass of each isotope by its natural abundance, expressed as a decimal, and then add the products.
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