The Building Blocks of Matter Chapter 3 - Atoms: The Building Blocks of Matter

From Philosophy to Science There were two schools of thought of the composition of the cosmos… is everything in the universe continuous and infinitely divisible Or, is there a limit to how small you can get? Particle theory was not the most popular early opinion, but was supported as early as Democritus in ancient Greece. From Philosophy to Science

From Philosophy to Science Democritus proposed that all the matter is composed of tiny particles called “Atomos” These “particles” were thought to be indivisible Aristotle did not accept Democritus’ atom, he was of the “matter is continuous” philosophy Because of Aristotle’s popularity his theory was adopted as the standard From Philosophy to Science

From Philosophy to Science By the 1700’s nearly all chemists had accepted the modern definition of an element as a particle that is indivisible It was also understood at that time that elements combine to form compounds that are different in their properties than the elements that composed them However, these understandings were based on observations not empirical evidence From Philosophy to Science

From Philosophy to Science There was controversy as to whether elements always combine in the same proportion when forming a particular compound. In the 1790’s, chemistry was revolutionized by a new emphasis on quantitative analysis because of new and improved balances This new technology led to the discovery of some new scientific understandings From Philosophy to Science

From Philosophy to Science The Law of Conservation of Mass: Proposed by Antoine Lavoisier States that mass is neither created nor destroyed during ordinary chemical reactions or physical changes. Which means the total mass of the reactants must equal the total mass of the products. From Philosophy to Science

Law of Conservation of Mass +  Carbon, C Oxygen, O Carbon Monoxide, CO Mass x Mass y Mass x + Mass y Carbon Monoxide, CO Mass x + Mass y + Carbon, C Oxygen, O Mass x Mass y 

From Philosophy to Science The Law of Definite Proportions: The fact that a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or the source of the compound NaCl is NaCl no matter if it is table salt (small crystals) or rock salt (large crystals) From Philosophy to Science

From Philosophy to Science The Law of Multiple Proportions: If 2 or more different compounds are composed of the same 2 elements, then the ratio of the masses of the 2nd element combined with a certain mass of the 1st element is always a ratio of small whole numbers From Philosophy to Science

Law of Multiple Proportions + = Carbon Oxygen Carbon Monoxide, 1:1 1 1 + Carbon Oxygen = Carbon Dioxide, 1:2 1 2

In 1808, John Dalton proposed an explanation for each of the proposed laws He reasoned that elements were composed of atoms & that only whole #’s of atoms can combine to form compounds His ideas are now called the Atomic Theory of Matter Atomic Theory

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, & other properties. ELEMENT 1 ELEMENT 2 ELEMENT 3 ELEMENT 4 Atomic Theory

Atomic Theory Atoms cannot be subdivided, created, or destroyed atoms of different elements combine in simple whole # ratios to form chem compds in chemical rxns, atoms are combined, separated, or rearranged + + Atomic Theory

Through these statements, evidence could be gathered to confirm or discount its claims Not all of Dalton’s claims held up to the scrutiny of experimentation Atoms CAN be divided into even smaller particles Not every atom of an element has an identical mass Atomic Theory

Atomic Theory Dalton’s Atomic Theory of Matter has been modified. What remains is… All matter is composed of atoms Atoms of any one element differ in properties from atoms of another element One of the disputed statements of Dalton was that atoms are indivisible Atomic Theory

In the 1800’s it was determined that atoms are actually composed of several basic types of smaller particles it’s the number and arrangement of these particles that determine the atom’s chemical properties. The def. of an atom that emerged was, the smallest particle of an element that retains the chemical properties of that original element. Atomic Theory

All atoms consist of 2 regions that contain the subatomic particles The nucleus The electron cloud around the nucleus The nucleus is a very small region located near the center of the atom In every atom the nucleus contains at least 1 proton, which is positively charged particle and usually contains 1 or more neutral particles called neutrons Atomic Structure

The electron cloud is the region that surrounds the nucleus This region contains 1 or more elec-trons, which are negatively charged subatomic particles The volume of the electron cloud is much larger than the nucleus Atomic Structure

Discovery of the Electron The discovery of the first subatomic particle took place in the late 1800’s. A power source was attached to two metal ends of an evacuated glass tube, called a cathode ray tube. A beam of “light” appears between the two electrodes called a cathode ray. Discovery of the Electron

cathode ray cathode anode Electric Current cathode ray cathode anode Cathode Ray

Modern Application: TV Picture Tube

Discovery of the Electron Investigators began to study the ray and they observed that… An object placed in the path of the ray cast a shadow on the glass A paddle wheel placed in the path of the cathode ray began to spin Cathode rays were deflected by a magnetic field The rays were deflected away from a negatively charged object Discovery of the Electron

Discovery of the Electron The first 2 observations support the idea that the ray is composed of tiny individual particles traveling through the vacuum tube The second set of observations support the evidence that the ray is composed of a substance that is negatively charged. Discovery of the Electron

Discovery of the Electron J.J. Thomson studied the rays and proved that they were tiny negative particles being emitted from the metal atoms. Dubbed these tiny particles “electrons” And it was later determined that the electrons were not part of the mass of the atom. Discovery of the Electron

What can their work help us conclude about the atom? atoms are composed of smaller particles, and one of these compo- nents is negatively charged atoms are neutral, so there must be an opposing (+) charge because electrons are essentially mass-less, an opposing substance makes up the mass of the atom First Atomic Model

Discovery of the Proton In 1886, E. Goldstein observed in the cathode-ray tube a new set of rays traveling in the opposite direction than the cathode rays The new rays were called canal rays and they proved to be positively charged And the particles mass were about 2000 X’s that of the electron Discovery of the Proton

Discovery of the Neutron In 1932, the English physicist James Chadwick discovered yet another subatomic particle. the neutron is electrically neutral It’s mass is nearly equal to the proton Therefore the subatomic particles are the electron, proton, and neutron. Discovery of the Neutron

Scientists still didn’t really understand how the particles were put together in an atom. This was a difficult question to resolve, given how tiny atoms are. Most thought it likely that the atom resembled Thomson’s model Atomic Structure

Negative particles embedded in a sphere of positive plasma-like matter. THINK… Chocolate Chip Cookie Mg + -

In 1911, Ernest Rutherford et al In 1911, Ernest Rutherford et al. provided a more detailed picture of the internal structure of the atom In his experiment, Rutherford directed a narrow beam of alpha particles at a very thin sheet of gold foil. Alpha particles (a) are He atoms that have been stripped of their electrons Rutherford Model

According to Thomson’s model, the heavy, positive alpha particles should pass easily through the gold, with only a slight deflection And mostly that’s how it happened. However, they found 1 in every 8000 particles had actually been deflected back toward the source. Rutherford Model

Rutherford suggested a new structural model of the atom. He stated that all the positive charge and the mass is concentrated in a small core in the center of the atom, AKA nucleus And that the atom is mostly empty space with electrons surrounding the positively charged nucleus like planets around the sun. Rutherford Model

Rutherford Model

With the exception of Hydrogen, every nucleus contains 2 kinds of particles protons and neutrons they make up the mass of the atom (Mass Number = Protons + Neutrons) Atoms are neutral because they contain equal #’s of protons & electrons Atomic Structure

The atoms of different elements differ in the # of protons in their nuclei and therefore in their positive charge The # of protons the atom contains determines the atom’s identity, also known as atomic number. Only Oxygen contains 8 protons Only Fluorine contains 9 protons Structure of the Atom

The nucleus is composed of a densely packed cluster of protons, which are all electrically positive Don’t like charges repel? Why don’t they fly apart? When 2 protons are in very close proximity, there is a strong force of attraction between them. similar attraction exists when neutrons are close Structure of the Atom

These short-range p+-n0, p+-p+, & n0-n0 forces hold the nuclear particles together, A.K.A strong nuclear forces. When these nuclear forces are strong enough the atom is stable If the forces are not strong enough the atom (heavier atoms) the atom is unstable and becomes radioactive. Structure of the Atom

Ch 3.3: Atomic Number

# protons in an atom = # electrons Elements are identified by the number of PROTONS they contain The “atomic number” of an element is the number of protons in the nucleus PROTONS IDENTIFIES AN ELEMENT!!! # protons in an atom = # electrons Because atoms are neutral!

Complete Symbol X Mass number Atomic Subscript → Superscript →

Cl 35 17 # OF PROTONS + # OF NEUTRONS MASS NUMBER ATOMIC NUMBER NUMBER OF PROTONS # OF PROTONS + # OF NEUTRONS

Mass Number Mass number is the number of protons and neutrons in the nucleus of an isotope: Mass # = p+ + n0 Element p+ n0 e- Mass # Oxygen - 10 - 33 42 - 31 15 Mass number 18 8 8 18 Arsenic 75 33 75 Phosphorus 16 15 31

Find the # of e-, p+ and n0 for sodium. (mass # = 23) Practice Problems Find the # of e-, p+ and n0 for sodium. (mass # = 23) Find the # of e-, p+ and n0 for uranium. (mass # = 238) Atomic # = 11 = # e- = # p+ # neutrons = 23-11 = 12 Atomic # = 92 = # e- = # p+ # neutrons = 238-92 = 146

Check for understanding: If an element has 91 protons and 140 neutrons find the: Atomic number Mass number number of electrons element name 91 231 91 protactinium

(Isotopes: The # after the name is the mass #.) An isotope refers to atoms that have the same # of protons, but a different number of neutrons. Because of this, they have different mass #’s. Examples---> (1) Carbon-12 & Carbon-13 (2) Chlorine-35 & Chlorine-37 (Isotopes: The # after the name is the mass #.)

Cl Cl 18 20 35 37 17 17 EXAMPLE OF AN ISOTOPE ATOMIC MASS ATOMIC NUMBER 18 NEUTRONS 20 NEUTRONS

Br Question #1 80 35 Find each of these: Atomic number Mass Number number of protons number of neutrons number of electrons Br 35

Question #2 If an element has an atomic number of 34 and a mass number of 78, what is the: number of protons number of neutrons number of electrons complete symbol

Units = atomic mass unit (amu) 12 Units = atomic mass unit (amu) The atomic masses listed in the Periodic Table are a “weighted average” of all the isotopes of the element.

In chemistry, chlorine has 2 isotopes: Weighted Average Practice Problems: In chemistry, chlorine has 2 isotopes: Cl-35 (75.8% abundance) Cl-37 (24.23 % abundance) What is the weighted average atomic mass of chlorine? 35 x 0.758 = 26.53 amu 37 x 0.2423 = 8.9651amu Add them up!!! (2) Oxygen has 3 isotopes: O-16 (99.76%) O-17 (0.037%) O-18 (0.2%) Estimate oxygen’s average atomic mass. Barely over 16.0 amu. + 35.4951 amu

Relating Mass Numbers to Atoms The Mole: the amount of a substance that contains as many particles as there are atoms in exactly 12 grams of carbon-12. Avogadro’s Number: the number of particles in exactly one mole of a pure substance = 6.022 x 1023. Molar Mass: the mass of one mole of a pure substance. Units = g/mol

This is when we get to use dimensional analysis! The conversion factors we need are: and of course…molar mass

Gram to Mole Conversions Mass of Element in Grams Number of Moles of Element Number of Atoms of Element

Practice Problem ALWAYS USE PARANTHESES AROUND YOUR CONVERSION FACTORS!! You have 3.50 mol of Copper. What is it mass in grams?