Chapter 2: Matter is Made of Atoms

Chapter 2: Matter is Made of Atoms Section 2.1 Atoms and Their Structures

Objectives Relate historical experiments to the development of the atom, Illustrate the modern model of an atom, Interpret the information available in an element block of the periodic table

Hypotheses, Theories and Laws Solving a problem: 1) Observation - use senses to observe the behavior of matter 2)HYPOTHESIS– testable prediction Hypotheses that are verified by repeating experiments

Hypotheses, Theories and Laws EXPERIMENT: Investigation (with a control) designed to test a hypothesis Hypotheses lead to scientific theories THEORY: Explanation based on many observations and supported by the results of many experiments. Ex: Dalton’s Atomic Theory

Hypotheses, Theories and Laws SCIENTIFIC LAW: Fact of nature that is observed so often that it becomes accepted as truth. A law can be used to make predictions, but does not explain why something happens Example: Sun rises in the east Theories can explain laws

SCIENTIFIC METHOD: (p. 57, Figure 2.5)

Early Ideas about Matter Greek Philosophers (2500 years ago) 4 Fundamental Elements- air, earth, fire, and water Questioned if matter could be divided endlessly into smaller pieces or if there was an ultimate small particle of matter Aluminum Foil

Democritus (dem-ock-rit-is) 460-370 B.C. Proposed the world is made up of empty space and the smallest particles of matter are called ATOMS This introduced the atomic theory of matter Different types of atoms exist for every type of matter

Modern Atomic Theory Antoine Lavoisier (Luh-voh-zee-ay) 1782 Concluded that when a reaction occurs, matter is neither created nor destroyed Law of Conservation of Matter

Conservation of matter and recycling Atoms are neither created nor destroyed You can’t throw anything away because there is no “away” Recycling Nitrogen- Nitrogen is atmosphere is converted into compounds used on Earth, then returned to atmosphere (p. 53, Figure 2.4) Recycling plastic, aluminum and glass- reusing atoms in these materials, we imitate nature and conserve natural resources (in natural processes atoms are recycled)

Joseph Proust 1799 Observed that the elements that composed compounds were always in a certain proportion by mass- LAW OF DEFINITE PROPORTIONS Ex: water is always 11% H and 89% O by mass

Law of definite Proportions

Dalton’s Atomic Theory of Matter (1803) Theory essentially intact with small modifications to accommodate new discoveries Main Points of Dalton’s Atomic Theory 1. All matter is made up of atoms 2. Atoms are indestructible and cannot be divided into smaller particles. 3. All atoms of one element are exactly alike, but atoms are different for different elements Late 19th century, experiments began to suggest that atoms are made up of even smaller particles (electrons, protons, neutrons)

Atomic Structure Today we know that atoms are made of smaller particles and Atoms of the same element can be nearly the same (but not exactly) Cathode Ray Tube Experiment- J.J. Thomson (1897) Vacuum tube- positive and negative electrode- ray travels through tube from the – to the +, rays bent toward + and away from the – Electrons, Protons, Neutrons ELECTRON- negatively charged subatomic particle Mass equal to 1/1837 the mass of a Hydrogen atom PROTON- positively charged subatomic particle The amount of charge on an electron and a proton is equal and opposite The mass of a proton is much greater than the mass of an electron (slightly less then a Hydrogen atom)

JJ Thomson Model

Atomic Structure Until 1910, it seemed that atoms were made up of equal numbers of electrons and protons J.J. Thomson discovered that Neon consisted of atoms of two different masses

Atomic Structure Ernest Rutherford- 1909 “Gold Foil Experiment” (p.62, Fig 2.9) + charged particles (alpha particles) sent through a gold foil – most went straight through (empty space), few deflected (hit nucleus) Revealed the arrangement of the atom- 1911 Atom is nearly all empty space with a small, dense, positively charged core called a NUCLEUS http://micro.magnet.fsu.edu/electromag/java/rutherford/

Rutherford Model

Atomic Structure Atoms of an element that are chemically alike but different in mass are called ISOTOPES The discovery of isotopes – atoms must contain a third type of particle that explains mass differences NEUTRON- neutral subatomic particle Mass is equal to that of a proton but has no electrical charge Existence of the neutron was confirmed in early 1930s

Bohr Model and Electron Cloud

Evolution of the atom (drawings)

Even smaller particles… Quarks – small particles of matter that make up protons and neutrons. 6 “flavors” or types – top, bottom, charm, strange, up and down An arrangement of 3 of these will form a proton, another arrangement will form a neutron.

Neutrinos Subatomic particles produced by the decay of radioactive elements Elementary particles that lack an electric charge F. Reines would say, "...the most tiny quantity of reality ever imagined by a human being

Neutrinos… Copiously produced in high-energy collisions Traveling essentially at the speed of light Unaffected by magnetic fields neutrinos Their unique advantage arises from a fundamental property: they are affected only by the weakest of nature's forces (but for gravity) and are therefore essentially unabsorbed as they travel cosmological distances

Atomic number and masses ATOMIC NUMBER: Number of protons in the nucleus of an atom of that element # of protons determines the identity of an element and many of its chemical and physical properties Atomic number also tells us the number of electrons in a neutral atom of an element (p+ = e-)

Atomic Number and Masses MASS NUMBER: Sum of the protons and neutrons in the nucleus of an atom Mass # = p+ + n0 n0 = Mass # - p+ Example: # of neutrons: mass # = 19 (Flourine) n0 = 19 - 9 = 10

Atomic Number and Masses Isotopes have different mass numbers because they have different numbers of neutrons (they have the same atomic number) Isotopes are identified by placing the mass number after the name or symbol of the element Ex: Li-7, Li-6 Ne-20, Ne-21, Ne-22

Periodic Table information Each box contains: element, state, atomic number, symbol and average atomic mass (weighted average of the naturally occurring isotopes) Atomic mass example: Chlorine has 2 isotopes- Cl-37 and Cl-35 24.2% is Cl-37 and 75.8% is Cl-35 So the atomic mass is 35.45

Protons, Neutrons, Electrons Particle Symbol Charge Mass in grams Mass in u Proton p+ 1+ 1.67 x 10-24 1.01 Neutron no 0 1.67 x 10-24 1.01 Electron e- 1- 9.11 x 10-28 0.00055 u – ATOMIC MASS UNIT (devised mass unit) 1 u = 1/12 the mass of a carbon-12 atom 1 u  mass of single proton or neutron

Chapter 2: Matter is Made of Atoms