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Chemical Names and Formulas. Chapter 5 And a little of chapter 14. We have just finished chapter 12, electron configurations. Why on Earth did we study it? That’s right, kiddos! To help us understand this chapter better… at least that’s the plan.

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chemical names and formulas

Chemical Names and Formulas

Chapter 5

And a little of chapter 14

slide2
We have just finished chapter 12, electron configurations. Why on Earth did we study it?
  • That’s right, kiddos! To help us understand this chapter better… at least that’s the plan.
  • To sum up Chapter 12 for our purpose today… we can predict where electrons should be in an atom of any element. We’ll use these predictions to determine if and how different atoms form compounds.
valence electrons
Valence Electrons
  • So far, when we’ve taken a look at electron configurations, we have predicted and accounted for every single electron in the atom.
  • But now we’re only going to focus on those electrons in the outermost energy level.
  • Why? Because those are the electrons that participate in the bonding ritual between atoms. Basically they’ll contribute to the stability and “happiness” of the atom. Trust me… it will make sense soon.
  • For example, let’s look at Oxygen. It has how many total electrons?
  • 8, that’s right! But how many electrons in its outermost energy level?
  • Only 6. How did I figure this out?
  • What is the highest energy level for oxygen? 2. It fills the s orbital completely and fills the p orbital with 4 electrons.
valence electrons continued
Valence Electrons continued
  • Why do we care if it only has 6? Well… oxygen wants to have a stable electron configuration like the closest noble gas, which in this case is Neon. How many electrons does Neon have in its outermost energy level?
  • 8. Aahhhh. Hmmmm….. Basically Oxygen wants to have all of its orbitals full like Neon. Right now it has 2 unpaired electrons which need buddies. So how can oxygen get to be stable and happy like Neon?
  • You bet! It needs to get 2 electrons from some other atom that needs to give some away.
  • What atom might need to give away 2 electrons? (Hint: look at the outermost energy level and determine which atoms have only 2 electrons there. )
  • Yep! Magnesium, Calcium… anything with only 2 electrons in the outermost s orbital.
  • Okay, so let’s make some predictions about our valence electrons for our representative elements.
  • Representative elements are group A elements 1A-8A. Basically these are the groups that would exist if we were to remove the transition metals from our P.T.
atoms vs ions
Atoms vs. Ions
  • We’ve now predicted the valence electrons for the representative elements.
  • If atoms gain or lose electrons, what happens to the overall charge of the atom?
  • It becomes positive or negative. When this occurs, the atom is now called an ion (charged particle).
  • Ions can either be positive (cation) or negative (anion).
  • Cations lose electrons to become positively charged. THESE ARE YOUR METALS!
  • Anions gain electrons to become negatively charged. THESE ARE YOUR NONMETALS!
  • There are a few exceptions (just like any other rule in science) but for the most part, this will be true. For example, Hydrogen is a nonmetal but will lose an electron to become positively charged. And there are polyatomic ions which are groups of atoms bonded together that carry a charge (usually negative).
slide6
Ions
  • Naming ions is simple:
  • Cation: same name as the element + ion
    • Mg+2 magnesium ion
    • Al+3 aluminum ion
  • Anion: change the element ending to ide
    • F- fluoride ion
    • S-2 sulfide ion
  • Pretty easy, right? We’ll practice more later.
slide7
Okay.. So we’ve talked about electrons, atoms, and ions, but not yet defined how they come together to form compounds.
  • Remember from chapter 1 that a compound is two or more elements that are chemically combined.
  • Unfortunately there is more vocabulary to include when discussing compounds:
    • Molecule: an electrically neutral group of atoms that act as a unit.
      • (Nonmetals combined with other nonmetals)
    • There are a few elements (7 of them) that only exist as molecules rather than individual atoms. They are called diatomic elements. Sometimes referred to as Super 7.
    • They are Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine, Bromine, Iodine.
    • H2, N2, O2, F2, Br2, Cl2, I2
    • These 7 elements when on their own will always come with a subscript of 2. DON’T FORGET THESE!
slide8
Molecular compounds: a compound composed of molecules.
    • These compounds have low melting and boiling points and are USUALLY gas or liquid at room temperature.
    • (Water and sugar are two examples here)
    • nonmetal + nonmetal
  • Ionic compounds: electrically neutral compounds composed of positive and negative ions.
    • These compounds have a crystalline structure, are solid at room temperature, and have high melting and boiling points.
    • (Table salt for example)
    • metal + nonmetal
chemical formulas
Chemical Formulas
  • Now we know what compounds are, at least a fuzzy picture but how do we write formulas for them?
  • Well… first we have to determine whether the compound will be an ionic (metal + nonmetal) or molecular (nonmetals only)
  • Let’s start with ionic compounds…
  • Binary ionic compounds are composed of two and only two elements. (Bi = 2) Remember that the overall charge on any compound must be equal to zero! That means we must have equal numbers of positive and negative charges.
ionic compounds
Ionic compounds
  • Let’s say I wanted to form a compound between lithium and fluorine. A cation and an anion.
  • First I need to figure out what charge each of the ions would have.
  • Li+ and F-  LiF, lithium fluoride
  • Because the charges equal each other, all I had to do was smack them together. … Okay its really not that simple on the atomic level but you get the picture.
binary ionic compounds con t
Binary ionic compounds con’t
  • But what if I try to smack together magnesium and chlorine? Do my charges equal out? Let’s see….
  • Mg+2 and Cl-
  • Since we learned before that we can’t have fractions of charges, how can I put these two ions together?
  • That’s right my little cherubs! I need 2 chlorines to balance out the magnesium so my result is….
  • MgCl2  This is called a subscript. It tells how many atoms are in the compound (if there’s more than 1).
  • Remember the overall charge must be zero, so my positive charges must equal my negative charges. ALWAYS ALWAYS ALWAYS!
  • Great! So, now that you’re all pros we can move right on…
  • How about another example…
  • Combine aluminum ion and sulfide ion.
  • Al+3 and S-2 Yikes! Now what? Remember no fractions or partial charges allowed!
binary ionic compounds con t1
Binary ionic compounds con’t
  • Al2S3
  • So, how did you figure that out? The overall charge of the compound must be zero so I needed to have 6+ and 6- in order to have a neutral compound.
  • Huh, you say? Try this instead… when you see the charges on two ions are unequal and you can’t just smack them together, then use the Travolta (or criss-cross) method.
  • Just remember to drop the + or – and make the number a subscript for the other element in your compound.
  • Okay.. Just watch me!
other ionic compounds
Other ionic compounds
  • Now you’re experts with binary ionic compounds, well.. almost
  • What happens if our compound contains one of those tricky (confusing, annoying, irritating,… I’m sure you’ll have your own name for it) transition metals.
  • But Miss I, what’s the big deal? They can still form binary compounds.
  • Yes, this is true. However, sometimes these funny elements can have multiple oxidation numbers (charges). For example, Iron can either have a +2 or +3 charge.
  • How do we tell the difference? Well… using roman numerals in the name or “reverse Travolta” the compound to figure out the original charge.
  • So, let’s do an example. Iron (III) oxide
  • Fe+3 and O-2 Travolta and viola!
  • Fe2O3
  • Alright.. How about copper (I) and nitrogen
  • Cu+ and N-3 Travolta… balance charges
  • Cu3N
  • No problem, right? Great. Let’s move on.
and more
And more
  • Not all ionic compounds are made of only two elements. Some contain a polyatomic ion.
  • What the heck is that? A polyatomic ion is a tightly bound group of atoms that behave as a unit and carry a charge.
  • Treat polyatomic ions like any other element. Just be careful when balancing charges and adding subscripts. Use parentheses to keep the polyatomic ion together and add the subscript on the outside.
  • Here are some examples of polyatomic ions:
  • Sulfate SO4-2
  • Nitrate NO3-
  • Carbonate CO3-2
  • Hydroxide OH-
  • Ammonium NH4+ Notice this one is positive!
and more1
And more
  • Let’s try a few:
  • Magnesium and sulfate
    • MgSO4
  • Potassium and hydroxide
    • KOH
  • Calcium and nitrate (be careful)
    • Ca(NO3)2
  • Aluminum sulfate (be careful)
    • Al2(SO4)3
naming ionic compounds
Naming ionic compounds
  • Basically, just say the name of the two parts…MgCl2 becomes magnesium chloride. We don’t need to say anything about the subscripts, just name the elements.
  • Same goes for ionic compounds with polyatomic ions. Just name the two parts… Al(OH)3. Aluminum hydroxide.
  • Just don’t forget the roman numeral with your transition metals. CuSO4. Copper (II) sulfate
  • See, its easy!
  • Let’s practice!
molecular compounds
Molecular Compounds
  • After spending so much time with ionic compounds, these should be a piece of cake!
  • Binary molecular compounds will contain how many elements?
  • 2 of course!
  • The big difference here is that there are no charges to balance because all of the elements involved are nonmetals.
  • So, how do they form compounds? Well, unlike when they’re bonding with metals and taking electrons from the metals, they actually have to share! We will talk about how they do this later on (next semester most likely).
  • All we’re doing now is learning how to name these compounds.
molecular compounds1
Molecular compounds
  • Molecular compounds differ from ionic in another way also. These compounds can combine differently.
  • Ever heard of carbon monoxide? What about carbon dioxide? Any idea what the difference is between them, and no its not that one will kill you. (They both can)
  • Yes, the prefix in front of oxide. That little prefix tells how many oxygens have combined with one carbon.
  • To write the formula:
  • Carbon monoxide: CO
  • Carbon dioxide: CO2
  • What about diphosphorus pentoxide?
  • Remember the prefix in front tells how many atoms there are. How many phosphorus? How many oxygen?
  • P2O5
  • You guys are genius I tell ya!
naming compounds and writing molecular formulas
Naming compounds and writing molecular formulas
  • Basically you need to:
    • Use a prefix to tell how many of each atom
    • The ending of the second element will always end in –ide
    • The vowel at the end of the prefix is dropped if the element starts with a vowel (like oxide. Monooxide just sounds funny)
    • Prefixes are ONLY for molecular compounds. DO NOT use them on ionic compounds.
    • If there is only one of the first element, “mono” is typically not used.
  • So there you go. More knowledge for your growing brain.
acids
Acids
  • No, you won’t be playing with these today. Bummer, I know. Besides, they’re really not that much fun.
  • Acids are compounds that donate a H+ ion when dissolved in water
  • Acids always begin with “H”
  • Binary acids have how many elements?
  • 2 What is always one of the two elements?
  • Hydrogen. My goodness you’ve learned quickly!
  • Unfortunately for you all acids are not binary. Some contain polyatomic ions. These will have different names and formulas than the binary will have. Big surprise I’m sure.
acids1
Acids
  • Binary acids: hydro_________ acid
    • HCl: hydrochloric acid
    • HBr: hydrobromic acid
  • No “O” hydro, hydro no “O”
  • Other acids:
  • Memorize these:
    • Sulfuric acid H2SO4 sulfate ion
    • Nitric acid HNO3 nitrate ion
    • Carbonic acid H2CO3 carbonate ion
    • Acetic acid HC2H3O2 acetate ion
    • Phosphoric acid H3PO4 phosphate ion
  • I ate the acid and got sick.
  • Last nite I was nauseous.