Atoms: The Building Blocks of Matter Chapter 3

OBJECTIVES Part 1: History of Discoveries How the structure of the atom was discovered The three fundamental laws of chemistry The Atomic Theory of Matter Part 2: Atomic Structure Counting protons, neutrons & electrons Isotopes, isotope notation Part 3: Weighing Atoms The atomic mass unit Average atomic mass Part 4: Counting Atoms The mole concept Converting between atoms, moles, and grams

The Particle Theory of Matter Chapter 3 – Section 1: The Atom: From Philosophical Idea to Scientific Theory The Particle Theory of Matter In 400 B.C. Democritus, a Greek philosopher, first proposed the idea of a basic particle of matter that could not be divided any further. He called this particle the atom, based on the Greek word atomos meaning indivisible. This early theory was not backed up by experimental evidence and was ignored by the scientific community for nearly 2000 years.

Foundations of Atomic Theory Three Fundamental Laws of Chemistry Chapter 3 – Section 1: The Atom: From Philosophical Idea to Scientific Theory Foundations of Atomic Theory Three Fundamental Laws of Chemistry By the late 1700s, experiments with chemical reactions led to the discovery of 3 basic laws: The Law of Conservation of Mass – Mass is neither created nor destroyed during ordinary chemical reactions or physical changes.

Foundations of Atomic Theory Three Fundamental Laws of Chemistry Chapter 3 – Section 1: The Atom: From Philosophical Idea to Scientific Theory Foundations of Atomic Theory Three Fundamental Laws of Chemistry The Law of Definite Proportions (Constant Compostion) – 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.

Foundations of Atomic Theory Three Fundamental Laws of Chemistry Chapter 3 – Section 1: The Atom: From Philosophical Idea to Scientific Theory Foundations of Atomic Theory Three Fundamental Laws of Chemistry The Law of Multiple Proportions – If different compounds are composed of the same two elements, then the ratio of the masses of the elements is always a ratio of small whole numbers.

Chapter 3 – Section 2: The Structure of the Atom Dalton’s Atomic Model An atom is the smallest particle of an element that has all the properties of that element. Atoms are too small to see…even through the most powerful microscope!! Dalton thought atoms were solid balls of matter and were indivisible.

Dalton’s Atomic Theory Chapter 3 – Section 1: The Atom: From Philosophical Idea to Scientific Theory Dalton’s Atomic Theory In 1808, John Dalton proposed an explanation for the three laws. His atomic theory states: All matter is composed of atoms. Atoms of the same element are identical; atoms of different elements are different. Atoms cannot be subdivided, created, or destroyed. Atoms of different elements combine in simple whole- number ratios to form chemical compounds. In chemical reactions, atoms are combined, separated, or rearranged.

Corrections to Dalton’s Theory Chapter 3 – Section 1: The Atom: From Philosophical Idea to Scientific Theory Corrections to Dalton’s Theory Dalton turned Democritus’s idea into a scientific theory that could be tested by experiment. But not all aspects of Dalton’s theory have proven to be correct. We now know that: Atoms are divisible into even smaller particles. A given element can have atoms with different masses.

Discovery of the Electron Chapter 3 – Section 2: The Structure of the Atom Discovery of the Electron In 1897, Joseph John (JJ) Thomson showed that cathode rays are composed of identical negatively charged particles, which were named electrons. The electron was the first subatomic particle to be discovered.

Thomson’s Cathode Ray Tube Experiment Chapter 3 – Section 2: The Structure of the Atom Thomson’s Cathode Ray Tube Experiment

Charge and Mass of the Electron Chapter 3 – Section 2: The Structure of the Atom Charge and Mass of the Electron In 1909, Robert Millikan measured the charge on the electron during his oil drop experiment. Using the charge-to-mass ratio, scientists were able to figure out the mass of the electron: about 1/2000 the mass of a hydrogen atom.

Millikan’s Oil Drop Experiment Chapter 3 – Section 2: The Structure of the Atom Millikan’s Oil Drop Experiment

Thomson’s Plum Pudding Model Chapter 3 – Section 2: The Structure of the Atom Thomson’s Plum Pudding Model After the work of Thomson and Millikan, the accepted model of the atom was called the plum pudding model. The atom was viewed as a ball of positively- charged material with tiny negatively-charged electrons spread evenly throughout.

Discovery of the Atomic Nucleus Chapter 3 – Section 2: The Structure of the Atom Discovery of the Atomic Nucleus More detail of the atom’s structure was provided in 1911 by Ernest Rutherford and his associates Hans Geiger and Ernest Marsden. The results of their gold foil experiment led to the discovery of a very densely packed bundle of matter with a positive electric charge. Rutherford called this positive bundle of matter the nucleus.

The Gold Foil Experiment Chapter 3 – Section 2: The Structure of the Atom The Gold Foil Experiment

Rutherford’s Atomic Model Chapter 3 – Section 2: The Structure of the Atom Rutherford’s Atomic Model After Rutherford’s gold foil experiment, the accepted model of the atom looked like this: A small, positively-charged nucleus with negative electrons surrounding it at some distance away. Most of the atom is empty space.

The Planetary Model of the Atom Developed when Niels Bohr used Rutherford’s model to determine that electrons could circle the nucleus only in certain “orbits” which we now call ENERGY LEVELS.

Chapter 3 – Section 2: The Structure of the Atom Subatomic Particles The nucleus is made up of at least one positively charged particle called a proton and usually one or more neutral particles called neutrons. Protons, neutrons, and electrons are often referred to as subatomic particles.

ATOMIC ACCOUNTING USING THE PERIODIC TABLE & ISOTOPE NOTATION

Chapter 3 – Section 3: Counting Atoms Atomic Number The atomic number (Z) of an element is the number of protons of each atom of that element. Atoms of the same element all have the same number of protons.

Chapter 3 – Section 3: Counting Atoms Mass Number The mass number is the total number of protons and neutrons in the nucleus of an atom. Atoms of the same element can have different mass numbers.

Chapter 3 – Section 3: Counting Atoms Isotopes Isotopes are atoms of the same element that have different masses. Isotopes have the same number of protons and electrons but different numbers of neutrons. Most of the elements consist of mixtures of isotopes.

Chapter 3 – Section 3: Counting Atoms Designating Isotopes Hyphen notation: The mass number is written with a hyphen after the name of the element. uranium-235 Nuclear symbol: The superscript indicates the mass number and the subscript indicates the atomic number. Mass number Atomic number

mass number − atomic number = number of neutrons Chapter 3 – Section 3: Counting Atoms Calculating Neutrons The number of neutrons is found by subtracting the atomic number from the mass number. mass number − atomic number = number of neutrons Nuclide is a general term for a specific isotope of an element.

Calculating Subatomic Particles Sample Problem Chapter 3 – Section 3: Counting Atoms Calculating Subatomic Particles Sample Problem How many protons, electrons, and neutrons are there in an atom of chlorine-37? Solution: Number of protons Number of electrons Number of neutrons = atomic number (on periodic table) 17 = number of protons 17 = mass number - protons 20

Calculating Subatomic Particles Sample Problem Chapter 3 – Section 3: Counting Atoms Calculating Subatomic Particles Sample Problem How many protons, electrons, and neutrons are there in an atom of Solution: Number of protons Number of electrons Number of neutrons = atomic number (on periodic table) 92 = number of protons 92 = mass number - protons 143

Writing nuclear symbols Sample Problem Write the complete nuclear symbol for the nuclide with 55 protons, 55 electrons and 78 neutrons. Solution:

Relative Units of Measurement  

Chapter 3 – Section 3: Counting Atoms The Atomic Mass Unit The standard used by scientists to compare units of atomic mass is the carbon-12 atom. One atomic mass unit, or 1 amu, is exactly 1/12 the mass of a carbon-12 atom. The atomic mass of any atom is determined by comparing it with the mass of the carbon-12 atom.

Chapter 3 – Section 3: Counting Atoms Average Atomic Mass Average atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element. The average atomic mass of an element depends on both the mass and the relative abundance of each of the element’s isotopes.

Calculating Average Atomic Mass Sample Problem 1 Chapter 3 – Section 3: Counting Atoms Calculating Average Atomic Mass Sample Problem 1 Copper consists of 69.15% copper-63, with an atomic mass of 62.93 amu, & 30.85% copper-65, with an atomic mass of 64.93 amu. What is the Average Atomic Mass of Copper? Solution Multiply the atomic mass of each isotope by its relative abundance. Add up all of the products. Divide by 100 and write amu(μ). Percent Abundance Mass Cu-63 69.15 62.93 Cu-65 30.85 64.93 x = 4352 x = 2003 + 6355 100 = 63.55 μ

Calculating Average Atomic Mass Sample Problem 2 Chapter 3 – Section 3: Counting Atoms Calculating Average Atomic Mass Sample Problem 2 A student believed that she had discovered a new element and named it Mythium. Analysis found it contained two isotopes. The composition of the isotopes was 19.9% of atomic mass 10.013 and 80.1% of atomic mass 11.009. What is the average atomic mass? Is Mythium was a new element? Solution: Average Atomic Mass: (19.9 x 10.013) + (80.1 x 11.009) 199 + 882 = 1081/100 = 10.81amu Because the atomic mass is the same as the atomic mass of boron, mythium was not a new element.

Chapter 3 – Section 3: Counting Atoms The Mole 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. It is a counting unit, similar to a dozen. In a dozen, there are 12 things. In a mole, there are 6.02 x 1023 “particles”. Visual Concept

ALL COMBINE IN THE SAME RATIOS The Mole Relates the atomic mass unit to grams, a unit we can work with! Reasoning: 1 atom of carbon-12 12.01 amu 1 mole of carbon-12 atoms 12.01 grams SINGLE ATOMS ALL COMBINE IN THE SAME RATIOS MOLES OF ATOMS GRAMS OF ATOMS

Chapter 3 – Section 3: Counting Atoms Avogadro’s Number Avogadro’s number: 6.02 × 1023 the number of particles in exactly one mole of a pure substance. Named for nineteenth-century Italian scientist Amedeo Avogadro. Visual Concept

Molar Mass  

Molar Mass  

Gram⟶Mole Conversions Chapter 3 – Section 3: Counting Atoms Gram⟶Mole Conversions Chemists use molar mass as a conversion factor in chemical calculations. Remember, molar mass means grams per mole. Example: What is the mass of 2.5 moles of Helium gas? Use molar mass as a conversion factor This is what we’re given: 4.00 g He 2.5 mol He x = 10. g He 1 mol He

Gram ⟶ Mole Conversions Sample Problem 1 Chapter 3 – Section 3: Counting Atoms Gram ⟶ Mole Conversions Sample Problem 1 What is the mass in grams of 3.50 mol of the element copper, Cu? Solution: Conversion factor Given 63.55 g Cu 3.5 mol Cu x = 220 g Cu 1 mol Cu

Gram ⟶ Mole Conversions Sample Problem 2 Chapter 3 – Section 3: Counting Atoms Gram ⟶ Mole Conversions Sample Problem 2 A chemist produced 11.9 g of aluminum, Al. How many moles of aluminum were produced? Solution: Conversion factor Given 1 mol Al 11.9 g Al x = 0.441 mol Al 26.98 g Al

Conversions with Avogadro’s Number Chapter 3 – Section 3: Counting Atoms Conversions with Avogadro’s Number Avogadro’s number can be used as a conversion factor between atoms and moles. Avogadro’s number units are atoms per mole. Example: How many moles of silver, Ag, are in 3.01  1023 atoms of silver? Conversion factor Given 1 mol Ag 3.01 x 1023 atoms Ag x = 0.500 mol Ag 6.02 x 1023 atoms Ag

Conversions with Avogadro’s Number Sample Problem 1 Chapter 3 – Section 3: Counting Atoms Conversions with Avogadro’s Number Sample Problem 1 What is the mass in grams of 1.20  108 atoms of copper, Cu? Solution: 2nd Conversion factor 1st Conversion factor Given 1 mol Cu 63.55 g Cu 1.20 x 108 atoms Cu x x 6.02 x 1023 atoms Cu 1 mol Cu = 1.27 x 10-14 g Cu