Unit 1

1. Unit 1 Elements and Atoms

2. Matter • Matter is the physical material of the universe. • We define matter by investigating its properties. • A property of matter is anything that we can use to distinguish one type of matter from another. • All matter is composed of one or more of over 100 elementary substances called elements. • Elements can not be broken down into simpler substances. • The atom is the smallest unit of mater that can still be defined as being one of the elements. • Molecules and compounds are formed when two or more atoms are bonded together chemically.

3. Pure Substnaces • A pure substance is matter that has distinct properties and a composition that does not vary from sample to sample. • A mixture is a combination of two or more pure substances. • There are two types of mixtures • Heterogeneous: • Where the individual particles of pure substances can been seen. • These types of mixtures are fairly easy to separate into pure substances. • Homogeneous: • Where the individual particles can not be seen. • This type of mixture is difficult to separate into pure substances.

4. Law of Definite Proportions • It has been observed that the elemental composition of any pure substance is always the same. • This means that the compound Sodium Chloride is always composed of one sodium atom and one chlorine atom: • NaCl • And that the compound calcium chloride is always composed of one atom of calcium and two atoms of chlorine • CaCl2

5. The Law of Multiple Proportions • Some elements can combine in different ratios to form different compounds with different properties. • Example: • Copper (I) Chloride – CuCl • Copper (II) Chloride – CuCl2

6. Percent Mass • The Law of definite proportions allows us to analyze compounds in terms of percent mass. • In any sample of NaCl 39.34 % of the mass will be sodium and 60.66 % of the mass will be chlorine. • Knowing the percent composition of a compound allows us to determine its empirical formula. • A compounds empirical formula is the simplest whole number ratio of the elements in the compound.

7. Hydrates • A hydrate is an ionic compound that is found in complex with water molecules. • Example: • The compound copper (II) chloride dihydrate has the formula… • CuCl2 * 2 H2O • Hydrates can be analyzed using very simple laboratory techniques.

8. The Mole • The mole is a counting number used in chemistry. • One mole of molecules is equal to 6.02 x 1023 of those molecules, just like one dozen eggs is equal to 12 eggs. • Talking about chemical reactions in terms of how many molecules are involved is not practice. • Talking about the mass of one molecule is not practice either. • The mole allows us to talk about a large number of molecules in an easy way.

9. Molar Mass • Counting molecules or atoms is not practical either. • Molar mass makes things easy. • The molar mass of a substance is the mass of one mole of that substance. • Example: • 1 mole, or 6.02 x 1023 atoms of sodium would have a mass of 22.98 g. • 1 mole, or 6.02 x 1023 molecules of water would have a mass of 18.02 g.

10. Molar Mass and Avogadro’s Number • We can use these two ideas to connect the mass of a sample and the number of particles in the same sample. • Example: • How many molecules would be in 100 g of LiBr?

11. Structure of the Atom • Atoms are the smallest unit of mass that can still be considered an element. • Atoms are composed of three basic subatomic particles… • Proton • Electron • Neutron • The protons and neutrons are found in the nucleus of the atom. • The electrons are found outside the nucleus.

12. Coulomb’s Law • Coulomb’s law describes the force of attraction between two charged particles. • F = k Q1Q2/d2 • Coulombs law tells us that as two charged particles get closer together the attractive force between them gets stronger. • This is very important when considering atoms and molecules.

13. Ionization Energy • The ionization of an atom or ion is the minimum amount of energy required to remove an electron. • The first ionization energy, I1, is the energy needed to remove the least tightly held electron from an atom. • Example: • Na (g) Na+(g) + e- • 495 kJ/mol • The greater the ionization energy the more difficult it is to remove an electron. • We can use coulombs law to justify trends in ionization energy. • Ionization energy is affected by two factors: • The distance between the electron and the nucleus. • The effective nuclear charge of the atom.

14. The further an electron is from the nucleus the lower the ionization energy will be. • This is because the attractive effect of the nucleus is less. • The more protons an atoms has the more positively charged its nucleus will be. • This will result in the electrons being attracted to the nucleus stronger. • Resulting in a higher ionization energy.

15. The Photoelectric Effect • It was observed that when light shines on an object electrons are ejected from the surface of the material. • This is called the photoelectric effect. • All light carries energy • E = hv • h = • v is the frequency of the light.

16. Photoelectron Spectroscopy (PES) • It was observed that not all light is able to eject electrons from all substances. • Only light of certain frequencies (i.e. certain energies) was able to eject electrons from different substances. • This led scientists to theorize that not all electrons are the same. • Scientists theorized that electrons occupied different energy levels that they call “Shells” and “Subshells”

17. Representations of Orbitals • The observation that electrons occupy different energy levels prompted scientists to come up with the idea of electron “Shells” and “Subshells”. • While the movement and location of electrons around the nucleus of an atom is quite complicated, this theory allows us to understand electrons and the energy associated with them better.

18. Energy Levels • The energy levels, or “Shells” that electrons can occupy around an atom are simply given whole number integers… • 1, 2, 3, 4 and so on • Each “Shell” contains the same number of “subshells” as the integer assigned to it. • For example energy level 1 contains only one subshell, while energy level 2 contains two and so on.

19. Subshells • To avoid confusion the subshells in each energy level are not assigned integers as names. • They are assigned letters. • The first subshell is assigned the letter s. • So the only subshell in energy level 1 is the 1s subshell. • Energy level 2 also has an ssubshell, called 2s. • It also has a second which is called 2p. • The third subshell (which is only filled in energy level 3 or higher) is d. • An the fourth is called the fsubshell.

20. Electrons and Subshells • All four subshells, s, p, d, and f, have a specific number of electrons that they are able to hold. • s can hold 2 • p can hold 6 • d can hold 10 • f can hold 14 • Just like the principle energy levels the subshells increase in energy. (s being lowest energy and f being highest.)

21. Electron Configurations • An atoms electron configuration is the way in which the electrons are distributed among the various energy levels and subshells. • The first thing we need to know when doing electron configurations is how many electrons an atom has. • Consider a neutrally charged atom of Lithium. • In order for it to have a zero charge it must have the same number of electrons as protons, 3.

22. Every element must have electrons that are in the very first energy level and subshell… i.e. 1s • However since this is an ssubshell it can only hold two of Lithium’s three electrons. • And since energy level one can only have one subshell we need to move to energy level two. • We start energy level to with subshell 2s • Where Lithiums 3rd and final electron would be. • We represent the idea that lithium has two electrons in the 1s subshell and one electron in the 2s subshell with the notation… • 1s22s1

23. Electron Configuration and Ionization Energy • We have already talked a little bit about ionization energy. • We have seen that as electrons are removed it gets harder and harder to do so. • We can use electron configurations and Coulomb’s Law to explain this trend. • Remember Coulombs Law is F = k Q1Q2/d2 • We saw already that elements with more protons in the nucleus had higher ionization energies. • But why is it harder to remove successive electrons from the same atom?

24. s • In an atom with many electrons the attractive force between the nucleus and the electrons in the outer-most subshell is low. • This is because these electrons are further from the nucleus. • But more importantly it is because the attractive force of the nucleus is lessened by the inner most electrons. • This is why it is easier to remove electrons that are further from the nucleus. • The electrons that occupy the highest, outer-most subshell are called valance electrons.

25. The Modern Periodic Table • Today’s periodic table is arranged based on electron configuration. • While there are some exceptions we can deduce an elements electron configuration based on it’s position in the periodic table.

26. Atomic Trends • One of the most important atomic trends is atomic radius. • In general: • Within each group (column) of the periodic table the radii of the atoms increases from top to bottom. • This is because the elements lower on the periodic table have more electrons. • within each period (Row) of the periodic table atomic radii decrease from left to right. • This is due to an increase in effective nuclear charge of the atoms as we move from left to right. • This draws the valance electrons in closer to the nucleus.

27. Ionic Radii • Just like atomic radius the ionic radius of an element is based on effective nuclear charge, and how many electrons is has. • Cations (+) are formed when electrons are removed from an atom. • Therefore cations are smaller than their parent atom. • Anions (-) are formed when electrons are added to an atom. • Therefore anions are larger than their parent atom. • For ions carrying the same charge, size increases as we move down the column in the periodic table.

28. Isoelectronic Atoms • Isoelectronic atoms are ones that contain the same number of electrons. • Examples: • O2-, F-, Na+, Mg2+, Al2+ • All of these ions have 10 electrons. • Each of these ions has more protons that the one before it. • This increases the effective nuclear charge of the ion. Drawing the electrons closer to the nucleus, and decreasing the radius

29. Trends In Ionization Energy • We already know that it is easier to remove the first electron from an atom than the second. • But what happens with we look at removing electrons from two different elements? • Within each period of the periodic table the first ionization energy (I1) increases from left to right. • This is because from left to right the atomic number increases, which increases the effective nuclear charge.

30. Within each group of the periodic table I1 decreases from top to bottom. • This is because the size of the atoms increases greatly. • Remember: • Atomic radii, ionic radii, and ionization energy all depend on effective nuclear charge and number of electrons. • Think about coulombs law • F = k Q1Q2/d2

31. Electronegativity • Electronegativity (AKA electron affinity) can also be studied using coulombs law. • Electronegativity is how strongly electrons are attracted to an atom to form an anion. • The most important factors when discussing electronegativity are effective nuclear charge and size. • Smaller atoms, with relatively high nuclear charges will have high electronegativites.

32. Quantum Mechanics • We have seen that Coulomb’s law is the basis for the describing the energy of interactions between protons and electrons. • Even though we call the space an electron occupies in an atom an “orbital” electrons do not follow specific orbits around the nucleus. • We have seen that only two electrons can occupy any given orbital and they must have opposite “Spin”. • Quantum mechanics uses advanced math to better describe the way electrons are arranged in an atom. • The Quantum mechanical model is still consistent with the electronic structures that correspond with the periodic table.

33. Atomic Models • The way scientists have thought about what atoms look like has changed many times. • One of the first scientists to come up with a useful atomic model was John Dalton. • During the period between 1803 and 1807 he came up with his atomic theory which postulates: • Each element is composed of extremely small particles called atoms. • All atoms of a given element are identical to one another, and different than atoms of different elements. • Atoms can not be created, destroyed, or converted into atoms of a different element. • Compounds are formed when atoms of more than one element combine.

34. Dalton Was Kinda Wrong • Mass spectroscopy is a tool used by modern day chemists to identify a substance. • It works by smashing molecules and atoms apart into smaller particles. • Mass Spec. experiments led to the discovery that atoms were not the smallest units of matter. • It also led to the discovery of isotopes.

35. Chemistry and Light • Light is more than just a beam. • Visible light is a type of electromagnetic radiation. (EMR) • EMR is composed of tiny particles called particles called photons. • The photons in a beam of EMR travel in a wave. • All photons carry energy associated with the frequency of that wave. • E = hv • Where v is the frequency of the light and his Planck’s constant (6.626 x 10-34 J-s). • Light used in chemistry is more often describe using wavelength. • Wavelength can be calculated using frequency and the equation: • v = c/λ • Where c is the speed of light 3.00 x 108m/s • The light we can see, and perceive as colors, ranges from 400nm to 800nm in wavelength.

36. Spectroscopy • Spectroscopy is defined as the study of the absorption and emission of light and other radiation by matter. • The type and amount of light that a molecule absorbs can tell us a lot about it. • One obvious difference in many compounds is color. • Molecules that are blue absorb and emit different wavelengths of light than molecules that are red. • Using the measured amount of light absorbed to determine something about a sample is called spectroscopy. • We will discuss the specifics of spectroscopy in Advanced Lab Class.

37. Conservation of Mass • In all chemical and physical processes the amount of mass before and after must be the same. • We can depict this idea through symbolic representations and particulate drawings. • In our next lab we will be investigation the following reaction. • 2 NaHCO3 Na2CO3 + CO2 + H2O

38. In our next lab we will use the idea of conservation of mass to determine the percent composition of a mixture of sodium bicarbonate NaHCO3 and sodium carbonate Na2CO3.